-------
                                   -14-
16% from the combustion of petroleum products, primarily residual


fuel oil.  The remaining emissions resulted from refining operations


(M%) and non-ferrous smelting (^10%) ,   A study on anthropogenic SO


production by the Study of Critical Environmental Problems (SCOP)


(M.I.T., 1970) concluded that globally 93 x 10  kg SO  were produced in


the year 1967-1968.  Kellogg et al. (1972) point out that this estimate


may be low because the "emission factor" used for this global estimate


is probably applicable only in the United States and will be higher


in other nations which rely more heavily on fossil fuels with high


sulfur contents such as coal,  Kellogg et al, believe an estimate of

              9
about 100 x 10  kg SO,, per year would be much more reasonable; of this,


almost 94% is emitted in the Northern Hemisphere alone.  Friend  (1973)

                                                        9
has estimated anthropogenic emissions of SO- at 130 x 10  kg per year.


      There are, of course, natural sources of sulfur dioxide as well


but the total amounts derived from these sources are extremely difficult


to quantify.  It is believed that most natural S0? is released by


volcanoes.  Kellogg et al. estimate that the quantity released by


volcanoes is about two orders of magnitude less than the amount  they

                                                            9
estimated to be a result of man's activities  (about 1.5 x 10" kg/yr


vs.-1-0 x 10   kg/yr).  There is, as the authors point out, a good deal


of uncertainty in their estimate.  Stoiber and Jepsen  (1973) estimated

                                              9
annual volcanic emissions of SCL to be 15 x 10  kg, an order of  magnitude


greater than the estimate of Kellogg et al„


      There are no other widely accepted natural S0~ sources,  Kellogg


et al„ have pointed out that despite the fact that SO  is so very


soluble  in water,  sea water might be a source rather than a sink if the

-------
right physicochemical conditions prevail.  Without  further  conv.i nc «>i:.




evidence though on this subject, one should probably  contend  thai-  : <.> ;




water would, in to to, provide only relatively minor amounts of  SO,,  11>




the atmosphere, if any at all






Removal Mechanisms




      Sulfur dioxide is very soluble in water; at 20°C  and  1  atmo;.j i'• •>




and either oxidizes to sulfate or photochemically reacts  with other




atmospheric contaminants.  Therefore, sulfur dioxide  is removed lrn;,i




atmosphere by various mechanisms involving water or other compound <;




The major identified sinks for this gas are:  precipitation scavrm-;"




chemical conversion, and absorption by soil, water, rock, and plan!




The lifetime of sulfur dioxide in the atmosphere is estimated to  i.i->r




between 20 minutes and 7 days (Nordo, 1973).  Following is  a  discus-.'




of the various mechanisms mentioned.






      Vegetation:  A portion of the much needed sulfur  used by  plan!-.




in metabolic processes has been shown by Fned(1948)  to be  atln ru>! >••!




to the direct absorption of S09 from the atmosphere,  especially in   •:




where the soil is sulfur deficient.  The ability of the plant to  r' >,''




this sulfur effectively without damage is dependent upon  the  rat'-  oi




absorption of S0? and the rate of production of sulffl-tes    The  [I'anl




can utilize this sulfur only if the production of sulfites  does uni




exceed the capacity of the plant to oxidize the latter  to sulfate1-




      Studies by de Cormis [1968) suggested that the  extent or  S00




absorption is directly proportional to the atmospheric  SO,,  cone en', ^ru

-------
                                    -16-
and is not influenced by the amount of sunlight.   Hill £1971) investigated


the uptake rates of several gases by an alfalfa canopy.   His results


confirmed those of de Cormis.  Hill found that SO  was absorbed with a


deposition velocity of 2.8 cm/sec, given a wind velocity of 1.8-2.2 m/sec,


as well as a number of other fixed variables.   He did note that, in general,


those gases readily absorbed by the alfalfa were those with the greatest


solubility in water [see Table II and Figure 1].





     Table II-  SOLUBILITY IN WATER AND UPTAKE RATE OF POLLUTANTS
Pollutant
CO
NO
°3
N02
S02
Uptake Rate
in Alfalfa*
2 9
[mol/m *sec) x. 10
0
2.1
34.7
39.6
59.0
Solubility
at 20°C
g/100 g
0,00234
0.00625
0,052
decomposes
10,8
 ''Concentration of the gas in the chamber was 2 x 10   mol/m  ,
      Using SO,, concentrations at 10 stations covering 5f>50 square


kilometers near Sudbury, Canada, as measured by Dreisinger and McGovern


(1970), and data from Figure 1, Hill calculated a hypothetical rate of


SO  removal by vegetation  (assuming a continuous cover of alfalfa) in


this area.  He indicates that the area is capable of removing 40 pg SO /

 2                3
m «sec or 560 x 10  kg/d,  His analyses make a number of assumptions

-------
                           -17-
O
                            234
                 Pollutant Concentration
                     (mol/m3)X I06

 Figure 1,  Uptake rates of different pollutants by an alfalfa

-------
                                   -18-
and yield a nice ball-park answer, but more important than the actual


value is the indication it gives as to how important vegetation can be


in cleansing our environment.


      Factors which influence pollutant uptake by plant canopies have


been discussed by Bennett and Hill (1973).  In a more recent communication


Bennett, Hill, and Gates (1973)  present a model simulating pollutant


transfer between leaves and the  free air surrounding it.  This model


is based on the rate of exchange via a series of external and internal


leaf mass transfer resistances„   The model indicates the importance of


gas solubility within the leaf.   Thus, use of deposition velocities


and vapor phase concentration alone, as done by various investigators


(Owers and Powell, 1974, Shepherd, 1974; Chamberlain, 1960; and Spedding,


1969) will be inadequate for the prediction of uptake rates of a gas


by plants.  In addition, variables which influence momentum and mass


transfer boundary layer thickness, such as wind velocity and gas


diffusjlvity, are shown to be important in determining the rate of


pollutant uptake by plants.  One should bear in mind from known mass


transfer theories that the shape of the plant and leaf would influence


the rate of uptake as well.  The effect of shape, however, has not


been tested or determined.


      Estimates of the amount of SCL removed each year by vegetal

                                                                   9
absorption vary greatly.  Eriksson's  (1963) cycle estimates 75 x 10  kg

                                                                 9
SO -S are removed in this manner.  Junge  (1963) estimates 70 x 10  kg


S02-S, Robinson and Robbins (1968) estimate 26 x 109 kg SO -S, Kellogg


et al.  (1972) estimate 15 x 10  kg S02-S, and Friend (1973) also

                      9
estimates that 15 x 10  kg SO -S are absorbed by vegetation each year.

-------
                                  -19-
Th e designation "SCL-S" is a means of expressing the amount of SCL
                   t-                                             i


present as its equivalent elemental sulfur weight.





      Soi1:   Recent studies have shown that soils are capable of absorbing



significant amounts of sulfur dioxide.  In 1936, Vandecave/e and associates



(as cited in Bohn, 1972) noted that acid  soils absorb only small



quantities of SCL,  Terraglio and Manganelli (1966), in studies of two



soil types,  found not only that SCL was more readily absorbed by soil



with a higher moisture content, but that the reaction also appeared



to be dependent upon the pH of the soil, more SO  being absorbed in



the soil where the pH was greater.  It is, of course, not that simple.



The degree of absorption is also dependent upon such factors as the



mineral and organic consent of the soil, soil structure, ion-exchange



capacity and porosity,  Faller (1968), as cited in Seim (1970), found



that the oxidation (a word Faller uses as synonymous with absorption)



of SCL was greatest in soils with the highest base saturation, other



factors being equal.  Faller and Herwig (1970)  suggest that the amount



absorbed is related to the exhangeable alkali and alkaline earth cation



content, and what is absorbed is converted to H«SO. when the water



content of the soil is sufficient (-20%).



      Experiments performed by Seim (1970) showed that soils can be



important media for absorbing SO =  He found that the amount of sulfur



absorbed by the 0-2 cm depth segment of a particular dry silt (Fayette



soil, analysis in his text) exposed to air containing 3720 pg SO /m


                                       2
for 24 hours was equivalent to 10.6 g/m ,   He does note that it is



questionable whether the soil can maintain this absorption rate for

-------
                                   -20-






extended periods of time,  but  it might be possible if sulfate is continually




removed from the soil by leaching and/or crop uptake,




      Smith et al ..  (1973)  showed that soils of varying pH,  organic




carbon content, and sand,  silt and clay content all readily absorbed




sulfur dioxide,  In a chamber containing air with an initial SCL




concentration of 270,000 yg/m  (100 ppm), all six soils reduced the




SO  content in the air by 95% in less than or equal to one  minute when




the soils were dry, and in less than one-half minute when the soils




were 50% saturated.  Sterilization of the soils caused no significant




change in their S0? uptake ability.  The removal mechanism  is




therefore probably not biological.  Like Seim, Smith et al.  suggest




that the SO  absorbed is oxidized to sulfate which may then be subject




to leaching and/or plant uptake.  If the sulfate content in the soils is




thereby reduced, the soils may maintain their SO  absorption capacity.




This hypothesis has not yst been investigated.




      Deposition velocities have often been used to determine the




removal of S07 from the air above soils and vegetation.  From the data




available in the literature, deposition velocities for soils appear to




be less than those for vegetation,




      Seim obtained deposition velocities for S07 of approximately 0.2




cm/sec for all the soils he investigated.  Chamberlain  (1960) reported




values of 0,3 and 0.7 cm/sec for average data over Great Britain.  For




combined soils and plants, Chamberlain used a value of 1.8 cm/sec.




Spedding (1969) calculated a maximum deposition velocity of 1.5 cm/sec




for plants.  Owers and Powell  (1974) calculated a mean deposition velocity




of 0,8 cm/sec over Great Britain assuming the countryside is all grassland.

-------
                                    -21-
Seiro points out that "if these values are valid and  the 0,2 cm/sec value


obtained for soils alone in this study should prove  to be valid over a


wide range of climatic conditions,  the conclusion which must be drawn is


that plants account for greater absorption of SO  from the atmosphere


than soil".


      Estimates on the amount of SO- that is absorbed by soils are


lacking and, for the most part, missing from most sulfur cycles that


have been compiled„  Unless this process has been taken into account


in estimating the total sulfur deposited by dry deposition on the land


surface then one must conclude that there is an obvious omission in

                                                   9
the cycles.  Eriksson (1963} estimated that 75 x 10   kg/yr of dry

                                                                 9
deposition sulfur goes through plants into the soil  while 25 x 10  kg/yr


are directly absorbed by the soil.   Abeles et al. (1971), based on


experiments they themselves ran, concluded that soils of the United

                                     9
States are capable of removing 4 x 10  kg of S0? per year,



      Rock:  Living matter has a threshold limit of  tolerance to pollu-


tants, below which no injury will occur, and, in fact, the substance might


actually benefit from the presence of that pollutant.  Rock reacts with


pollutants such as SO , at all concentrations.  The  results of these


reactions may not be visible for quite some time, for the effects are


cumulative.  The extent of breakdown of the rock is  therefore less a


product of the momentary concentration of the pollutant than the uptake


per unit time on a unit area of the material (Luckat, 1973).


      The effects of SO- on frescoes, monuments and  other edifices


have been most pronounced over the last century, expecially in Europe


where high-sulfur coal and oil are used as heating fuels.  The basic

-------
                                   -22-
destructive reaction is that of sulfuric acid (SO- + 1/2 CL + hLO -»•



1LSO.) on the carbonate matrix of limestone and sandstone in the presence



of moisture.  Spedding (1969b) has shown that as the relative humidity



in the air increases, the SO- uptake rate by oolitic limestone increases



significantly (Table III).
          Table III.  LPTAKE OF S02 BY OOLITIC LIMESTONE
Relative
Humidity
11
13
79
81
so2
concentration
Ug/m3
360
.280
LOO
370
Time of
exposure
20
40
48
10
Uptake
yg S02/cm2
of surface
0.069
0.061
0.24
0.28
Uptake
rate
5.0
2.2
7.2
40.3
      The product of the; reaction between sulfuric acid and the carbonate



matrix is gypsum if sufficient evaporation occurs:





                CaCO- + H SO, + H00 -*• CaSO/ 2H.O + C0_
                    3242        42      2




Because the calcium carbonate is slowly being replaced by gypsum the



rock would be subject to increased weathering rates due to:



       (1) an enhanced chemical disintegration caused by the much



          greater solubility of gypsum in water than calcium



          carbonate.  The newly formed mineral would be subject

-------
                                  -23-
          to dissolution in water with accompanied leaching out



          of the rock, and



      (2)  an enhanced physical disintegration caused by the almost



          two-fold volume expansion in the rock accompanying the



          formation of gypsum.



      Other properties such as the density and porosity of the rock



are also important and affect the amount of weathering to be expected



in a rock per unit uptake rate; a rough, porous, lime-cemented



sandstone would be expected to weather faster than a smooth, dense



limestone (Luckat, 1973).



      To determine whether the absorption of S07 by sedimentary rocks



would constitute a significant sink for SO , a number of assumptions



were made:



      (1)  Knowing that approximately 30% of the total earth's surface


                              14  2
          area of about 5 x 10   m  was land (Holmes, 1965), and



          assuming that perhaps 5% of the land surface had exposed



          rock (F. E, Wickman, 1974), and of this approximately 75%



          is sedimentary (Leet and Judson, 1965), it was calculated



          that about 1% of the total earth's surface was covered



          with rock capable of absorbing SO .



      (2)  Luckat (1973) and Spedding (1969b) found SO  absorption

                                       2
          rates ranging from 5-200 mg/m -d in sandstone and limestone,



          respectively.  Luckat's measurements were taken in a highly



          industrialized region of Germany;  Spedding's measurements



          were made with S0? concentrations approximately 100 times



          greater than the average world-wide background concentration,

-------
                                    -24-
                                                         2
          For this calculation,  the lower limit  of 5 mg/m -d


          was used,

                                                                   9
      From these values it  was calculated that approximately 9 x 10  kg

                  n
SO /yr or 4.5 x 10"' kg S/>r could be removed by stone under these


optimal and exaggerated conditions.  Comparison of this value with those


in Table V shows this value to be considerably smaller than that of any


other natural SO  sink.


      There are many problems related with making any such estimate as


was done above.  First of all, any estimate of the total percentage


of the earth's surface which has exposed rock is entirely speculative


at this point in time because geological maps are not available for all


parts of the world.  Then again, there is the problem of defining


an outcrop and mapping it in its strictest sense; that is, mapping the


outcrop without magnification and without the inclusion of soils or


detritus as part of the outcrop.  There is also the fact that not all


the sedimentary rock that is exposed is sandstone or limestone; much


of it is shale or mudstone or something of similar density which would


probably not absorb SCL to any extent.  Finally, the areas where most


of the rock does outcrop would be in areas where the S0_ level would


probably be quite low.  It can therefore be assumed that rock constitutes


a negligible sink for SO  on a global scale,




      Water Bodies:  Theoretical arguments in support of the contention


that sea water is capable of absorbing significant quantities of sulfur


dioxide from the atmosphere first appeared in a paper by Liss  (1971)


and were modified by Liss and Slater  (1974).  Briefly, the rate at which


a gas is exchanged across an interface can be obtained by using a mass

-------
                                  -25-
transfer coefficient,  k, having velocity dimensions and defined as:
                            k = flux of gas	
                                concentration difference            l }
                                over a thickness, z
      The overall mass transfer coefficient, K, based on liquid (£) or

gas (g)  phase concentrations is related to individual mass transfer

coefficients by the following equation
                     1=1+1      I  = 1    M.
                     K0   k0 + Hk  °r K    k  + k0                  ( J
                      £    £     g     g    g    £
                     R£ = r£ + rg or Rg = ^ + rg                   (3)
where k,, and k  are the mass transfer coefficients for the gas and
              o

liquid phases respectively.  The terms r,, and r  are measures of the
                                               o

resistances in the liquid and gas phases and are obtained from the

reciprocals of individual mass transfer coefficients.  The Henry's law

constant, H, is defined as:
     _ equilibrium concentration in gas phase (g/cm  air)	
       equilibrium concentration of un-ionized dissolved gas in       '
       liquid phase (g/cm^ H-O)
and differs for every gas.

      Whether the net transfer of a gas is controlled by the liquid

phase or gas phase is determined by the ratio of the resistances of

the gas to liquid phase (r /r,,) .   For most gases (N_0, CM. and CC1  tu

name just a few) r /rc < 1, and transfer is controlled by the liquid pha-
                  g  *

-------
                                   -26-
Suifur dioxide transfer though is evidently controlled by the gas phase;

that is r /r0»l.   For water vapor r. = 0 and the exchange across
         g  x,                       x,
the interface is controlled by processes in the gas phase.  With the

knowledge of mass transfer coefficients for one contaminant, one can

approximate the mass transfer coefficients for other contaminants.

      Liss and Slater (1974) used the mass transfer coefficient for

water vapor to calculate the overall mass transfer coefficients for

a number of gases crossing the air-sea interface.  The results are

given in the following table::
     Table IV.  MASS TRANSFER COEFFICIENT FOR A NUMBER OF GASES
                CROSSING THE AIR-SEA INTERFACE
k <
Gas g , H
cm/ sec cm/ sec
K*
g £ cm/sec
S02 0.45 9.6 3.8 x 10"2 573 0.45 (g)
NO 0.53 0.0055 1.6

CO 0.67 0.0055 50

CH, 0.885 0.0055 42
4
HO 0.833 «>
6.6 x 10 3 0.0055 (£)
-4
1.7 x 10 0.0055 (£)
-4
1,5 x 10 0.0055 (£)
0.83 (g)
 *The overall exchange constant, K, is expressed on either a gas  (g) or
 liquid  (£) phase basis.
       Liss  (1971) first showed that the exchange of S0? with aqueous

 solutions was a function of pH  (see Table V).  A chemical enhancement

 factor, a,  equals the ratio of the kfl values for a reactive and an
                                    A/

-------
                                   -27-
     Table V.  CALCULATION OF THE RATIO OF THE GAS TO THE LIQUID

               PHASE RESISTANCE
pH
2
2.8
3
4
5
6
7
8
9
a
2.7
11.7
18.0
169
1,376
2,884
2,966
2,967
2,967
k£(S02)
cm/sec
0.0075
0.0325
0,050
0.469
3,822
8.011
8,239
8.242
8.242
sec
130
30.
20.
2.
0.
0.
0.
0.
0.
cm
8
0
1
26
125
121
121
121
sec cm £
32 0, '
32 I,M
52 1 f.
32 15 1
32 123,1
32 25b, I)
32 264 L
32 26; ',
32 264,5
inert gas exchanging under identical conditions.  As can be seen, abov<



a pH of 28 the gas phase resistance becomes dominant, and therefore



the exchange of SO  with natural waters (pH 4-9) is controlled by 1 IK



gas phase.  At a pH of 8 (approximately that of sea water) the gas pi' "



resistance has increased to more than 99 percent of the total resistm^



Brimblecombe and Spedding (1972) experimentally confirmed these resiiHr



of LisSc  Liss and Slater (1974) suggest that the exchange of" othei



gases, such as NH , SO  and HC1 which also are very soluble and
                 O    »J


undergo rapid hydration reactions, might also be controlled by the

-------
                                    -28-
resistance of the gas phase,   It should be pointed out that the difference




in the values of the ratio r /rff for SCL at a pH of about 8 between
                            g  )6       z



Tables IV and V is due to (1) a smaller chemical enhancement factor used




in calculations for Table IV yielding a slightly larger k0 ,„  , value,

                                                             2

and (2) a difference in calculating the value of k .„ ,.  Whereas Liss




and Slater derived k ,   , by correcting the value of k  ,   ,  by a factor




equal to the ratio of the square roots of the molecular weight of water




to the molecular weight of SO  (as in Table IV), Liss (1971) had assumed




tnat K f,, 0^ -K  fr^r* -» «
      g(H20)  g(S02l


     Spedding (1972) collected samples of sea water and,  by measuring




the amount of SO  absorbed by the water from an atmosphere containing




SO , determined the deposition velocity of S0? over sea water.  He found




••"hat the deposition velocity of SO  increased linearly as the gas flow




rate increased.  The natural buffering capacity of sea water prevented




any drop in the pH  f-8), and thus would not limit the amount of SO




absorbed.  Earlier, Terraglio and Manganelli  (1967) had found a rapid




and substantial decrease in pH when SO  was absorbed by distilled water.



     Most data  on SO  solubility in the literature have been determined




when atmospheric S0? concentrations in the experimental chambers far




exceeded those  found in ambient air.  Hales and Sutter (1973) worked




towards closing this obvious gap by running a number of experiments




to help "quantify the relationships between S09 solubility, concentration,




and hydrogen-ion impurity at levels normally  encountered  in nature."




The dissolution of  S0~ was assumed to proceed according to the reactions




set forth by Falk and Giguere  (1958) :

-------
                                     -29-
                        S0?   + 2H20 + H30  + HSO~                     (6)

                          "aq






                        HSO~ + H20 J H30+ + S03~                       (7)






The second ionization (reaction 7) was assumed to be negligible.




     Based upon the first two  of the above reactions, Hales and Sutter




derived an "extrapolation equation, relating the concentration of total




dissolved SO- in water (c n ) to airborne concentration and solution
            2.            oU«


acidity" as follows:
                 [SO.]     -[H,0+]   + /[HJD+I2  + 4 K.  [S00] /H
                 1  2J_g_      L 3  Jex     L 5  Jex	1  L  2Jg
                   H
where those terms bracketed are concentrations in moles per  liter  and




[H 0 1  is the "excess" hydrogen ion concentration, defined  as the
  O  C.A.


concentration of hydrogen ion in solution present due to  sources other




than the dissolving of SO .  K  is the equilibrium constant  for reaction  ((>)




and H is the Henry's law constant.  Extrapolation of SCL  solubility  data




from Johnstone and Leppla (1934) down to low ambient SO   concentrations




show deviations ranging from 0.7 to 21.2% of those by Hales  and Sutter,




In general, the percent deviation increased as the concentration of  SO




in the gas phase decreased.  The authors suggest that even though  this




deviation does exist, the ability of equation  (8) to predict low SO-




concentration solubility appears excellent and they recommend its  use




whenever low concentration solubility data are needed.

-------
                                     -30-
     A more exact equation for determining the total dissolved SCL in

water can be derived based on the solution equilibria involved.  The

resulting equation not only takes into consideration the pH of the solution

and the atmospheric partial pressure of the gas, P   , but applies to
                                                  oU--i
different aqueous phase conditions, including both sea water and rain

water.  Also it eliminates the use of the often confusing and now outdated

hydronium ion (H,0 ) .   The dissolution and dissociation reactions for S09
                
-------
                                    -31-
where [SO . H,,0] is the activity of solvated sulfur dioxide and Pqn   is


the SO  partial pressure in atmospheres.  The defining equations for


the ionization constants are:
                              [H] [HSO~]


                         K  =
                          l   [S02-H20]
                              [H+] [SOI]

                         K  = - —                               (14)

                                [HS03]




Using equations (12),  (13) and  (14), a molal mass balance for the system



can be written
                               aq
wherein,                          „ p


                                     S02

                         mSO    = ~Y -
                          bU2      YS09
                             aq       2
                              n        aq
                                 Ki H Pso2

                         •HSO-, - 7— rr-T,
                             *   YHS03  L  J
                                 K2 Kl
                         m     =

                            3
     Substituting the values for the molality for each of these  species


into equation (15), the following general formula for the concentration  of
dissolved S09 is derived:
-------
                                     -32-
            2        2      2aq                           3




     The definition of pH is




                               pH = -log [H+]




Therefore, by convention, in equations the activity of the hydrogen ion


                 +      -r>H
can be written [H ] = 10 *'  as desired.



     For use in predicting SO,, solubility, equation (19) requires values



for H, K , K  and the three activity coefficients in addition to the



pH and the SO  partial pressure.   Scott and Hobbs (1967) give H = 1.24,



K  = 0 0127, and K0 = 6.24 x 10~  at 25°C,  Corresponding data at other



temperatures are given by Johnstone and Leppla (1934).   One case of



considerable interest is the absorption of S0~ in sea water.  Activity



coefficients of aqueous sulfite species, SO «H?0, HSO_, and S0~ in sea



water are not available.  However, Reardon (1974) determined activity



coefficients for the analogous carbonate species and also sulfate ion in



sea water.  From a point of view of similarity of species and coherence among



the variables, the sulfate activity coefficient is rejected in favor of



use of the carbonate species activity coefficients.  Hence,
                         YS02  S YH2CO°

                             aq
-------
                                     •O-S-
By substituting these values into equation (19),  and assuming that sea



water has a pH of 8.1, equation (13) reduces to
                        Lmcn  = 8,42 x 107 Pcn                         (20)
                          b(J                  2
where the amount of dissolved SO  is seen to be a linear function of the



atmospheric partial pressure of SO-.



     The minimum background concentration of S0_ is 1 yg/m  which



corresponds to 3.8 x 10    atmospheres SO  partial pressure.  From



equation (20) the amount of S09 dissolved in sea water at equilibrium


                 _2
would be 3,2 x 10   mol/JL  As an upper limit one might consider ambient



SO- concentrations founa by Luckat  (1973) in highly industrialized



sections of Germany.  The observed 360 yg/m  is equivalent to P    =


        -7                                                       2
1,4 x 10  atm.  In equilibrium with such an atmosphere sea water



would absorb 11.6 mol SO /£.  This molarity is an order of magnitude higher



than that where Henry's law is known to hold.  Thus, the predictive



equation might fail under these circumstances.



     When the dissolution of S02 takes place in rain water, the activity



coefficients required for equation  (19) may be assumed equal to unity



because the ionic strength of rain water (a measure of the interionic



effect resulting primarily from electrical attraction and repulsions



between the various ions) would be very low, probably on the order of


  -4
10  .  Therefore, equation (19) reduces to




                                            9            — 1 0

                                           '             "
                                     [H ]          [H ]2
-------
                                     -34-
ni this case the hydrogen ion activity is dependent upon the amount



of SO  absorbed and cannot be specified a priori.  The charge balance



for the dissolution and dissociation reactions of S02 in water is
                           = [HSO~] + 2[SO] + [OH"]
By substituting in the respective expressions for the [HSO ] and  [S0~]
                                                          o         o


as given by equations (12, 13, and 14), and setting  [OH~] = 10    /[H+],



an expression is derived for the partial pressure of S0  as a function of
  +
[H] as follows:
                                   [H+]3 -  10~14
                     p    =  ---
                                                    _
                                      -?  +             -q
                        2    1.58 x 10   [H ] + 1,97 x 10





Choosing an initial value for the [H ] , and substituting that value



into equation  (22), a corresponding value for Pcn   (atm) can be found.


                                                 2
These values for Pcn  and [H ] are then in turn substituted into
                  oUrt


equation (21) to find the amount of total dissolved SCL,



     In pure rain water the pH is 7.0,.  Any SO,, which dissolves would



produce an acid solution and a drop in pH.  Thus, for pH = 7 , P    =


                                                                  2      -9
Zmcn  = 0,.  For a pH of 5, PQO  and Sm    are calculated to be 6.3 x 10
  S02                       bU2       bU2


atm and 1.0 x 10   mol/£, respectively,,  For a pH of 3, P    = 6.35 x
                                                         ou«

  -5                         -5
10   atm and £mcn  = 1,1 x 10   mol/£.  Clearly as  Pcn  increases, the
               oU,                                  ou_
                 L.                                    <-

amount of dissolved SO  increases and the pH decreases.  The dissolved



S0? predicted by equations (21) and  (22) agree excellently with the



experimental data of Hales and Sutter  (1973) and Terraglio and



Manganelli  (1967),
-------
     Estimates of the amount of SO  absorbed annually by the oceans


indicate the importance of this sink,  Liss and Slater (1974) estimate


this SCL flux at 1,5 x 10   kg/yr based upon their own calculations.


Their estimate is in good agreement with those of Eriksson (1963)  (2 x


TO11 kg/yr) and Robinson and Robbins (1968) (0.5 x 1011 kg/yr), but is


lower than that determined by Spedding (1972)  (9.6 x 10   kg/yr).  Liss

                                                           3
and Slater explain that this discrepancy is due to a 3 yg/m  difference


in the mean atmospheric SCL concentration used by each and because


Spedding's value for the total resistance of the gas phase (1/K ) was
                                                               to

much lower.  As noted earlier, Kellogg et al.  (1972) consider the net


flux of S00 from air to sea to be negligible,  based on observations made


by Pate et al. (person^ communication to Kellogg et al.) that in some


areas, where the equilibrium vapor pressure of SCL in surface waters


exceeds the partial pressure of SCL in the air above it, the ocean


might actually be a source of SO .   Although this condition may exist


locally, on a global basis it would make far more sense to consider


the ocean as a sink in light of the high solubility of SCL in water.


Also, whereas laboratory experiments have shown that the solution and


rate of oxidation of S0_ to sulfate in distilled water is limited by


the pH as the solution becomes more acidic (Terraglio and Manganelli, 1967),


this would not be the case in the ocean (Spedding, 1972).  The ocean has


a natural buffer capacity that maintains the pH at about 8.1-  Therefore,


pH would not be a limiting condition in sea water and so it should have


the capacity to constantly absorb S02-   This seems consistent with the


discovery of Liss (1971) and Liss and Slater (1974) that it is the gas


phase and not the liquid phase that limits the transfer of S0? across


the air-sea interface.
-------
                                    -36-
     Washout and Rainour, of SO :   The major portion of SO  present in the

atmosphere is probably removed by the processes known as rainout and

washout.  Rainout invokes the scavenging of SCL and sulfate particles within

the clouds while washout, is the removal of these sulfur compounds below

cloud level via precipitation.  The SO  scavenged will undergo a series

of reactions, some catalytic, and ultimately form H~SO. drops or a

sulfate salt.  The following paragraphs, due to their nature, might just

as well be included in the section on chemical reactions involving S0?

in the atmosphere, but they are included here so the reader can best

follow the sequence of events from the time the gaseous SO  reacts with

water vapor in the atmosphere until the sulfate salts are removed by

precipitation or dry deposition.

     From the time SO,., is absorbed by cloud droplets it is both innized

and oxidized by reactions (9), (10) and "(11) shown in the previous

section and  (23)and (24) as follows  (Miller and de Pena, 1972):
                                      OH  = HSO~ + H20                (23)
                          S0~ + 1/2 02         + S0~                  (24)
                                      dissolved


It is obvious then that any SO  absorbed by water droplets will not only

contribute to the sulfate content of rainwater but will also cause a

decrease in the pH of the rain.  The pH of rain usually ranges from 5-6

but can drop down to about 4  in areas where large amounts of SO  are

present in the atmosphere.  Because the pH drops, reaction  (2s) is

important only in the first few seconds.  Perhaps it would be easier

to visualize if the sequence  above is given as the net reaction:
-------
                                    -37-
                     S0o + H00 + 1/2 0_         -»• SO. + 2H+          (25)
                       /    2.         Z , .    -.    ,    4
                                       dissolved





     Because the oxidation of SO,, in the liquid phase does not occur at a



rate fast enough to account for the sulfate content found in rain,



investigations were undertaken to find an effective catalyzing agent.



Experiments have shown that of all tne metals found in the atmosphere,



Mn salts were the most effective in promoting SO  oxidation (Junge and



Ryan, 1958; Johnstone and Coughanowr, 1958; and Matteson et al., 1969).



Although Mn salts are the best known catalyst for SCL oxidation, the concen-



tration of these salts in the atmosphere is still not great enough



to account for the sulfate content in rain.  Recent investigations though



have shown that when ammonia is present the rate of sulfate production in



solution is greatly enhanced (van den Heuval and Mason, 1963; Scott and




Hobbs, 1967; and Miller and de Pena, 1972),



     The results of Scott and Hobbs'initial study as to how the



concentrations of SCL and NH  effect the amount of S0~ produced and the



pH of the precipitation are shown in Table VI,  The results are given



in terms of the SO? concentration because the rate of SO. production



depends only on  [S0~]    (reaction 22)  The first line of Table VI gives



the initial values of the pH and the S0~ concentration corresponding



to partial pressures of SO  and NH  typical of those found in the



Earth's atmosphere, namely 20 ug/m   (7 x 10   atm) and 5 yg/m  (7 x 10



atm), respectively.  The following lines show the effect of varying the



concentrations of S02 and NH .  One can see that more sulfite is produced



given the same SO,, concentration when NH  is present.  Also, as the



concentration of NH  increases so does that of the sulfite produced.
-------
                                    -38-
     Table VI,,   INITIAL VALUES OF THE pH AND THE CONCENTRATION OF SULFITE

                (SO*) ION FOR VARIOUS PARTIAL PRESSURES OF SO  AND Nil..
                   o                                         2.       . t
Partial Partial
pressure pressure Initial
of S0? (atm) of NH (atm) pH
*•• O
1 x 10~9 '7 x 10~9 6.34
7 x 10~9 0 4.97
1.4 x 10"8 7 x 10~9 6.21
7 x 10"9 1,4 x 10~8 6»48
5 x 10"6 5. x 10"6 6,35
5 x 10"6 0 3.55
Init j al
concentration
of SO'
(moles liter )
3.4 x 10~5
6.0 x JO"8
3,, 6 x 10 "
6.2 x i(f J
2.4 x 10"2
6,2 x 10"8
One can also see the effect that the presence of NH  has on the pH,  When
                                                   o


no ammonia is present the pH is highly acidic.  When present though, the



hydroxyl ion produced upon the dissolution of ammonia in water tends to



neutralize the hydrogen icns produced by the dissolution of SO ,   Therefore
                                                              Zj


when ammonia is present the pH remains a little over 6,  These values



have all been attained with C0? present in the gas phase at an average



concentration of 362,000 yg/m  (311 ppm-).  Therefore the acid pH is a



result of dissolution of C0_ as well as SO?J but the variability is



mainly attributable to S02 and NH  concentrations,



     Field investigations by Beilke and Georgii (1968) indicated that



the absorption of gaseous S0? by rainout and washout accounted for '/5%



of the sulfate content in rain water and that scavenging of sulfate
-------
                                    -39-
paiticles contributed only 25%.  The model formulated by Miller  an-1




contradicts those measurements of Beilke and Georgii,  Miller  and  i-




show that the sulfate content of rainwater is much more dependent,  on




scavenging of particles rather than SCL.  In their model, the  huli.'i




content of rain water near a highly concentrated SO  plume of  or. i ;••




moderate particle concentration, showed that the contribution  of- ••<




the total sulfate concentration was 4 times less than that of  ;>ui( •!'




particles.  Miller and de Pena's model makes more sense, especJdM,




because all reactions involving S09 in the atmosphere ultimate!;-  i




to the formation of SO. .




     Appreciable effort has been devoted to the analysis of SO  ,  <




ging by rain (Engelmann, 1968; and Fuquay, 1970).  Field measureinu i




SO  washout (Hales, Thorpj and Wolf, 1971) have demonstrated that Un




a significant accumulation of SO  in the water drops, and consequs, >:i




calculations of the washout based on deposition velocities are jna-d




     A comprehensive analysis of reversible washout based on the




interaction of raindrops with atmospheric contaminants has been  PI-,




by Hales  (1972),  This analysis indicates the use of overall mass  <••




coefficients for determining the washout.  From this analysis, li  j




apparent that the degree of success in determining washout rat^s •!  •




on estimating mass transfer coefficients and solubility data,  J1;;J«'




used mass transfer coefficients for the limiting cases of gas  phu'j*




control or liquid phase control without internal circulation wJLJu1




the drop to obtain the upper and lower limits of gas washout raito.




Further study (Hales, Dana, and Wolf, 1973; and Hales, Wolf, and f);m •




1973) has led  to the development of a mathematical model for  jpr
-------
                                    -40-
ground level concentrations in the rain as a function of location beneath



a plume under stable meteorological conditions.





     Atmospheric Reactions Involving SO :   Reactions involving S09 in



the dry state, not unlike those discussed above for SCL in the wet state,



are very complex.  The most important reaction involved here is the



photochemical oxidation of SCL which takes place in polluted atmospheres.



     Early measurements of the rate of photo-oxidation of SO  made by



Gerhard and Johnstone (1955) are the most widely quoted.  They found the



rate of SCL oxidation tc proceed from 0.1-0.2%/hour,  Recently, Cox and



Penkett (1970) measured this oxidation rate in experiments using purified



ambient air and found rates ranging from 0.04-0,65%/hour.  Analysis of



the gaseous composition in their test chamber found small amounts of


                                   -4      3
hydrocarbon compounds (.04 -CL2x 10   mol/m , or, 0.1-0.5 ppm) &nd



nitrogen oxides  (<10 yg/m  , or, <0.005 ppm), and they suggest that their



presence might be responsible for the high observed photo-oxidation rates.



Renzetti and Doyle  (1960) had earlier suggested that the rate of photo-



chemical aerosol formation  (the end result of the photo-oxidation of SCL)



is greatly accelerated in the presence of olefinic hydrocarbons and



nitric oxide.  Endow, Doyle and Jones  (1963) and Harkins and Nicksic



 (1965) have shown that the resulting aerosols consist almost entirely



of sulfuric acid droplets when the relative humidity is greater than



or equal to 50%.



     Cox and Penkett  (1971a) have given experimental evidence to  support



the hypothesis that low  concentration olefinic hydrocarbons and nitric



oxide can greatly affect the rate of SCL photo-oxidation in air.
-------
                                   -41-
Results of their experiments can be easily observed in Figure 2.  The


rate constant obtained for the oxidation of SO  in this experiment was


0.025 hr'1 with 40 ug/m3 (0.03 ppm) NO and .04 x 10"4 mol/m3 (.1 ppm)


cis-2-pentene present.  They argue that due to reduced light intensity


in their chamber, their rate obtained would be analogous to a conversion


rate of about 10%/hour over London.  Although this conversion rate seems


very high, they suggest it gives support to the oxidation rate of SO


of 0.65%/hour found in their earlier experiment as being more realistic


for existing atmospheric conditions than those found by Gerhard and


Johnstone,


     Although the rate of S02 oxidation by ozone alone is quite slow,


Cox and Penkett (1971b) found that when SO  was injected into a chamber


containing ozone and olefins the oxidation rate was greatly enhanced.


They also found that the oxidation rate, or aerosol formation, was


dependent upon the nature of the olefin; the rate of aerosol production


was much slower for terminally unsaturated propene and 4-methyl-1-pentem


than it was for the internally unsaturated olefins, cis-2-pentene and

                                                                  3
2-methyl-2-butene.  They calculated the oxidation rate of 270 ug/m


(.1 ppm) S0? in the presence of 2 x 10   mol/m  (,05 ppm) ozone and


olefin to be approximately 3%/hr for cis-2-pentene and 0.4%/hr for proper-


     As little as is known about the oxidation of SO  in the troposphere,


still less is known about its oxidation in the stratosphere.  Cadle and


Powers  (1966) have suggested a possible 3-body reaction with atomic


oxygen



                           SO,, ••- 0 + M -* SO, -•• M
-------
                          -42-
  1000.s
ro
 E
 x
 a»
 c
 o
 c
 a>
 o
 c
 o
 o
CO

TJ

 O


"5
 «/}
 O
 i_
 (D
 O

 <3-
O
CO
 CVJ

T.
   100.
10.
       o S02 concentration

       A Aerosol concentration
                                 Inject

                          NO(.3xiO~4mol/m3)

                          2Pentene (.4l6X!0"4mol/m3)

                                 I
                   100        200

                      Time(min.)
                                        300
 Figure 2.  Aerosol formation and SO  decay during the photooxidation of SO ,
-------
                                    -43-
wherc M is a molecule of 0~ or N0, which acts to carry off excess  energy,



thereby preventing prompt reversal of this reaction.  Mulcahy, Steven,



and Ward (1967) found the rate constant for this reaction at room



temperature and from 0,7-3 torr  (0.9-4x10   atra) to average  (2.7±0,5)


    56    -2    -1
x 10  m  rnol   sec  „   Friend, Leifer and Trichon  (1973) calculated



the rate of SO  removal from the atmosphere by this reaction to proceed



quite slowly; 6 x 10" % hr~  at sea level and 2 5  x 10~4?6 hr"1  at 20 km.



     Davis, Payne, and Stief (1972) suggested that  the reaction of  SO...



with the hydroperoxyl radical might also be of importance in the strato-



sphere




                           S00 -i- HO,, ->• S07 + OH
                             2.     <-     j




They found the rate constant for this reaction to  equal 1,8 m  mol



sec   (t factor of 3).  Based on estimated stratospheric H09 levels



made by Nicolet (1972), Friend et  al. (1973) concluded that this



reaction would ultimately lead to  the production of comparable if  not



greater amounts of sulfate then would be produced  by the reaction  of



S09 with atomic oxygen



     Friend et al„ (1973) also offer a reaction sequence showing the



3-bodied hydrolysis of S0_ to H SO  and the subsequent hydrolysis  of the
                         O     £   *-T


sulfuric acid to H_SO. solutions,  or acid embryos.    These solutions



are neutralized by their reaction  with ammonia to  form "salt embryos".



The salt embryos act as nuclei for growing stratospheric particles on



which the oxidation of SO  occurs  as follows:
                     2SO  +  2H 0 + 0   embry°>  2H SO  (solution)
-------
                                    -44-
Th s NH  acts as a catalyst.   When the ammonia is depleted from the




surrounding atmosphere, Friend et al. suggest that further absorption




and oxidation of SCL on the particle will be inhibited due to a drop




in pH as the buffer capacity of the NH. ion is exceeded.   This reaction




sequence is thought to be responsible for the layer of HuSO. and sulfate




particles found in the lower stratosphere.




     Perhaps the best way to summarize the possible reactions involving




SO  in the atmosphere would be by repeating the summary made by Robinson




and Robbins (1968).  "It seems that in the daytime and at low humidity,




photochemical reaction systems involving SO , NO  and hydrocarbons are




of primary importance in t.ie transformation of SO  into essentially an




H SO. aerosol.  At night and under high humidity or fog conditions,




or during actual rain, it seems that a process involving the absorption




of SO  by alkaline water droplets and a reaction to form S0~ within




the drop is a well-documented process and can occur at an appreciable




rate to remove SO  from the atmosphere."






Environmental Sulfur Cycle




     Table VII summarizes the estimates made by Eriksson  (1960), Junge




(1963), Robinson and Robbins  (1968), Kellogg et al. (1972) and Friend




(1973) in compiling their respective sulfur cycles.  It should be kept




in mind that while all the values given are only estimates, the degree




of uncertainty in some is greater than that in others.  For instance,




there are no measurements upon which estimates of the quantity of H?S




or S~ emitted by decaying land and sea biota can be based.  Also,




there are no specifications of the form of sulfide emitted  (i.e.,




H2S, HS~, or S~).  These estimates have therefore been arrived at by the
-------
-45-






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nj
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-------
                                   -46-
authors' balancing of particular portions of their cycles; the difference




needed to balance each section has then been set equal to the flux of




il S or S  from the land and sea to the atmosphere, respectively.





                          GLOBAL SULFUR CYCLE





     A schematic global sulfur cycle is shown in Figure 3.  This




diagram is a slightly modified version of that devised by Friend  (1973).




The only modifications other than layout is the addition of atmospheric




oxidation schemes for H S and SCL, as well as a flux for the sorption of




SO  by soils (Eriksson, 1963).  The other fluxes shown in this diagram




are listed in column 5 of Table VII.  All fluxes related to the




atmospheric sulfur cycle are discussed separately in the sections




concerning the sources and sinks of these compounds.  The fluxes  between




the pedosphere, hydrosphere,and lithosphere have been discussed in detail




by Friend (1973).  All fluxes are given in Tg S/yr  (10   g S/yr).




     The sulfur inventory of the various reservoirs or spheres are




encircled and given in Tg S.  All inventories are as shown by Friend  (1973)




     From cycles such as this the residence time of sulfur, or the time




taken for a complete turnover of the burden in a reservoir, can be




calculated for both the atmosphere and hydrosphere.  The residence




time, T, of a substance is defined as T = M/R, where M is the burden  or




inventory of the substance in a reservoir, and R is the total flux in-




to  or  out of that reservoir.  The residence time of sulfur in the




atmosphere then is equal to






                T = M/R =1.8 Tg/(217 Tg/yr) = 0.0083 yr = 2.7 days
-------
i
!

|

|
S
!•
fr

°r
~)
J
1 \
! i








..^



&
J
!
I





j









1
i
i
i
E
i
i
i
1
J
l'il
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I- [
•:.". I
';! ^

i
{
>
i
j



s j
-------
                                   -48-
(Friend, 1973).   The residence time of particular sulfur compounds would
be different.   Hidy (1973) has calculated the atmospheric residence time
of SO  to be 4 days and that of H S to be 2 days.  Nordo (1973) has estimated
the residence time of SO  to range between 20 minutes and 7 days.  Individual
estimates of the residence time of sulfate in the atmosphere have not been
made, although it is probably comparable to that of sea salts, or about a
day or two (Eriksson, 1963).
     Using Friend's data a calculation of the residence time of SO. in the
oceans can be made as follows:


                 T , M/R =  .("±0.1) x 108Tg =  (4_22±,324)  X1Q6 years
                             3.08 x 10  Tg/yr

or about 4 million years.  This estimate is a bit under half that of
Mackenzie and Garrels (1966).  They calculated the residence time of S0~
in the oceans to be 10 million years based on river input data,  Holser
and Kaplan (1966) have calculated the residence time at 21 million years
based on their own geochemical cycle.  The values they use for the
sulfur  inventory of the ocean and fluxes out of the ocean are  essentially
the same as Friend's.  However, they do not include an estimate  of the
sulfur  removed by marine plants.  As a result, their fluxes are
underestimated by about 164 Tg/yr.  Even without this flux though, the
residence time they calculated was not reproducible.  The information given
in their cycle leads to a  SO  residence time of  about 9 million  years.
-------
                                    -49-
                             CARBON CONTAINING GASES



CARBON MONOXIDE


     Each year more carbon monoxide is released into the atmosphere than


any other pollutant (excluding carbon dioxide), and, each year the quantity


released increases.  One would therefore expect a gradual increase in


ambient CO levels,  yet one finds that the background concentration of


this gas in the atmosphere has not fluctuated the last few decades.


There must then be one or several major active sinks for CO within the


troposphere.  Until just a few years ago though, investigations on


possible sinks had only turned up additional sources of CO,  This dilemma,


as a result, came to be known as the "CO sink anomaly".

                                                           3
     Background concentrations of CO range from 47-230 yg/m  (0.04-0.20 ppra)


(Jaffe, 1973).  Robinson and Robbins (1968)  indicate that a mean concen-

                   3
tration of 100 yg/m  (0.1 ppm) is found in the northern hemisphere,


while concentrations less than 58 yg/m  (0.05 ppm) would probably be more


common in the southern hemisphere„



Sources


     Carbon monoxide is the product, of incomplete combustion of fossil


fuels containing carbon.  With the advent of large scale industrialization


and the tremendous increase in the use of the automobile, great quantities


of CO have been emitted into the atmosphere.  Recent investigations for


CO sinks to explain why the ambient CO concentration has not been increas-


ing have resulted in the discovery of new natural sources of CO whose total


quantity far exceeds the total mass of CO produced as a result of man's


technology (Stevens et al., 1972).
-------
                                   -50-
     By far the greatest single anthropogenic source of CO is motor


vehicle exhaust,  Jaffe (1973) estimated that of a total anthropogenic

                                                             q
CO emission source in the United States in 1970 of 132,6 x iO  kg, 96.9

    g
x 10  kg resulted from the burning of gasoline by motor /eludes alone.


Other significant contributions to this man-made CO burden are from

                              9
solid waste disposal (6.5 x 10  kg), industrial process loss (10.3 x

  9                                        9
10  kg) and agricultural burning (12.5 x 10  kg).


     Jaffe estimated the total CO emitted by combustion processes on a

                                                  9
global scale for 1970 to be approximately 360 x 10  kg,  lie notes that


whereas the level  of CO produced by man in the United Slates appears


to be leveling off, globally it is on an increase,  Und;jr.ievoloped


nations, which are undergoing increased technological development are


not, as yet, concerned with the resulting pollution as muL.h as the


economic gains and so their emission levels are on the n ,e


     The most widely recognized natural source of CO is forest fires

                                              9
which have been estimated as releasing 11 x 10  kg CO into the atmosphere


each year  (Robinson and Robbins, 1968),,  This is, by no means, the only


major natural source of CO.  Jaffe  (1973) has written an excellent


review on  all possible natural CO sources.  The following is condensed


from that paper,


     Minor amounts of CO have been found to be released from volcanoes


and marshes  (Flury and Zernik, 1931).  CO can also be formed during


electrical storms  (White, 1932) and by the photodissociation of CO


in the upper atmosphere  (Bates and Witherspoon, 1952); Calvert, Kerr,


Demerjian  and McQuigg, (1972) have suggested the photodissociation of


formaldehyde as a possible source of CO and recently, Swinnerton,


Lamontagne and Linnenbom  (1971) found CO to be present in rain water
-------
                                    -5J-



in rather high concentrations.  Galbally (1972) offered a hypothesis


wherein the CO in rain is a product of the photodecomposition of aldehydes


in the rain water by sunlight.


     The ocean was first suggested as a major source of CO by Swinnerton,


Linnenbom and Lamontagne (1970),   Linnenbom, Swinnerton and Lamontagne

                                                           q
(1973) have estimated the oceans  can produce up to 220 x 10  kg each year,

                                                                  q
whereas Liss and Slater (1974) have estimated this flux at 43 x 10  kg per


year.  Robinson and Moser (1971)  suggested that plants could indirectly

                              9
be the source of about 54 x 10  kg CO by the oxidation of released


terpenes.  Finally, McConnell, McElroy and Wofsy (1971) suggested that

                      9
approximately 900 x 10  kg CO are produced each year by the oxidation


of methane.


     In light of this new source  information, Stevens et al. (1972) belici",


that natural sources of carbon monoxide could yield about 10 times more


CO than all anthropogenic sources in the northern hemisphere.  Up to this


time it has been assumed that anthropogenic activity released far more


carbon monoxide than nature.  It  will be interesting to follow the out-


come of this contradiction over the next few years.



Removal Mechanisms


     Carbon monoxide can be regarded, for all intents and purposes, as


being insoluble in water; its actual solubility being only 0,00234g/100g


HO at 20°C and 1 atm.  Therefore, wet processes such as washout and


rainout can be regarded as playing an insignificant part in the removal


of CO from the atmosphere.  Experiments on the absorption rate of CO by


an alfalfa canopy (Hill, 1971) showed that virtually no CO was absorbed


and therefore, vegetation can be  disregarded as a sink.  Absorption of
-------
                                     -52-






CO by the oceans can now be disregarded as well because it has recently




been shown (Swinnerton et al.,  1970) that the oceans actually constitute




a significant natural source of carbon monoxide.  It seems then that




the major sinks for CO are gas-phase reactions in the troposphere and




stratosphere and absorption by soil fungi (Inman, Ingersoll and Levy,




1971, and Inman and Ingersoll,  1971).




     There are two schools of thought concerning the sources and removal




mechanisms of CO in the atmosphere. The old school believed that anthro-




pogenic emissions far surpass natural emissions of CO into the environ-




ment, and, due to lack of information at that time, did not recognize




that chemical reactions in the stratosphere and troposphere play an




important part in destroying CO.  Estimates of the residence time of




CO in the atmosphere in this school ranged from less than 4 years




(Bates and Witherspoon, 1952)  to 2.7 years (Robinson and Robbins, 1968).




The new school recognizes the importance of natural sources of CO to the




tropospheric inventory as well as the importance of both stratospheric




and tropospheric reactions in removing it.  Residence times estimated




by this school are on the order of 0.1-0,2 years  (Weinstock, 1969;




Levy, 1971; Dimitriades and Whisman, 1971; and Levy, 1973b).  In light




of the new information available on sources and sinks of CO, it may be




unfair to segregate residence time estimates into those arising from




old and new schools of thought.  For clarification purposes the separa-




tion does seem necessary.






     Soil:  Experiments by  Inman and Ingersoll  (1971) showed that both




potting soils and natural spoils absorbed significant amounts of CO.  They




observed that, in general,  soils with the highest uptake activity were
-------
                                    -53-
those with higher organic content and lower pH.   Inman et al.  (1971)  also


noted that if a soil  was autoclaved (sterilized)  removal of CO by the


soil was inhibited.  This suggested that the removal was due to biological


activity in the soil.  Table VIII (Inman et al.,  1971) summarizes the


results of tests on several natural soils performed in the laboratory.


Further tests isolated some aerobic micro-organisms which were present


in the soil and were then tested for their CO uptake abilities.  The


following fungi were found to be active in removing CO from the soil:


4 strains of Penicillium digitatum, 1 strain of Penicillium restrictum,


4 strains of aspergillus, 1 of Mucor hiemalis, 2  of Haplosporangium


parium and 2 of Mortierella vesiculato.


     In late 1972 Ingersoll published the results  of a more extensive


study on the uptake of CO by soils.  He measured  the in situ uptake at


various locations throughout North America.  Following is a short


summary of his findings.


     1)  Total amounts of CO destroyed by various  soils ranged from

                    2
7.5 to 109.0 mg CO/m «hr,the spectrum ranging from tropical soils which


were the most active down to desert soils which were the least active.


Although there were many exceptions, more CO tended to be destroyed by


soils with low pH and moderate moisture content (^20%) than others.


     2)  The rate of CO uptake decreased as the concentration of CO


decreased in the air, with maximum removal with ambient concentrations of


100,000 yg/m3 (100 ppm).


     3)  Given the same soil, CO uptake was far greater when vegetation


was growing than when the soil  was under cultivation.  Ingersoll


suggests this is because the amount of organic matter present  in soils
-------
-54-
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-------
                                    -55-
being cultivated is substantially lower than that present in soils



actively growing crops.



     4)  Soils that were removed from their site of origin and tc^'



in the laboratory showed greatly reduced uptake ability.  The magnu



of decrease was not uniform from one soil sample to the next^



     Based upon data he had collected which he corrected for tempo*  .



and uptake variations, Ingersoll (1972) estimated the CO uptake pu;



of the conterminous United States and the world.  His estimates oi


        9                 12
505 x 10  kg and 14.3 x 10   kg CO per year for each, respective]},


                                          3
based on ambient CO levels of 100,000 yg/m  (100 ppm), three ordei•



magnitude greater than average ambient CO levels.  Ingersoll no tec1



though that Seiler (197?) had found soil CO uptake rates to be one



of those Ingersoll had measured when these soils were exposed to atii:



CO levels (230-1200 yg/m ) rather than concentrations of 100,000 )ii;



On the basis of this information, Ingersoll reduced his estimates <•



total capacity of soils to consume CO in the United States and the.


                 9                12
to around 50 x 10  kg and 1.4 x 10   kg per year, respectively.  ft.



Seller's article has not been published, it is impossible to state



whether his uptake measurements were taken in the field or the lab.



Ingersoll, as mentioned above, has found that uptake rates for thr



soil were greatly reduced when taken in the laboratory as compared



field measurements.  Unless measurements were taken by both Ingerso•



and Seiler under the same conditions, one must question the valid!i



in using Seiler's data to reevaluate estimates made using Ingersol'



data.  Since the estimates made at this stage are only a first



approximation, perhaps such concerns are unimportant.  Earlier, In,
-------
                                    -56-
Ingersoll (1971)  also estimated that soils of the continental  United

                                             g
States were capable of removing over 500 x 10  kg CO per year.


     Based on measurements of the CO uptake rate of two soils, Heichel


(1973) concluded that the resistance of the soil to CO removal increases


as the moisture content of the soil decreases, and also increases as the


degree of air circulation above the soil decreases.  Comparing his


uptake and resistance measurements with those of Turner, Rich  and


Waggoner (1973) who measured the uptake rate of ozone on one of the same


soils, Heichel found the rate of CO removal by the soil to be  2-3 orders


of magnitude less than the removal of 0 .  There was also a correspond-


ingly greater resistance to removal of CO than 0 .  Heichel determined


on the basis of his mt-asurements, and assuming an average ambient CO


level of 350 yg/m  (0.3 ppm), that the land can remove the equivalent


of 14% of the tropospher:.c CO inventory of 4.75 x 10   kg CO  (Weinstock

                                           g
and Niki, 1972) each year, or about 67 x 10  kg/yr.


     Recent experiments by Smith et al. (1973) also showed that soils


are capable of effectively removing CO from the atmosphere.  They also


found though that when mcist soils were placed in the chamber with air


containing 100,000 yg CO/m  (100 ppm), the concentration of CO in the


chamber rose, in one case tripled, before it was reduced to zero.  This


effect was found to be more pronounced in sterilized soils than


unsterilized soils.  They surmised that the evolution of CO from the


soils is a nonbiological process, but made no attempts to identify the


processes responsible.  Seiler and Junge  (1970) also noticed this effect.


Their test results showed that at ambient CO levels (200 ug/m  } or 0.2 ppm)


the rate of CO production by the soil  is balanced by the rate of CO
-------
                                    -57-
absorption by the soil.  This finding lends another complication to the
identification of sources and sinks for CO.  Perhaps ultimately it will
be shown that soils do constitute a major sink for carbon monoxide; first
we need to know if there is evidence of CO evolution from in situ soils
or whether this is a phenomenon only found in laboratory experiments,
and we need quantitative data to see what the ratio of CO evolved to CO
removed from the atmosphere is at different ambient CO levels.

     Atmospheric Reactions Involving CO:  Recently it has become
apparent that the stratosphere constitutes a sink for carbon monoxide.
The significant reaction seems to be the reaction of CO with the hydroxyl
radical (Weinstock, 1959; Pressman and Warneck, 1970; and Pressman et al.,
1970) as follows:
                              OH + CO -> CO  + H                        (1)

The rate constant for this reaction at -56°C, the temperature of the
                                    4  3
lower stratosphere is k  = 6.02 x 10  m /mol*sec (Schofield, 1967).
Westenberg (1972) found that the rate constant increases to 10.2 x 10
 3
m /mol-sec at room temperature (27°C) .  Reaction 1 is followed by the
sequence:
                           H + 02 + M -> H02 + M                       (2)

                            CO + HO  -»- CO  + OH                       (3)
                                   L*     £.

The rate constant for reaction 2 has been found to be equal to 4.35 x
  O 7  A    9
10"  m /mol -sec (Westenberg and de Haas, 1972).
-------
                                   -58-
     The significance of reaction 3 has recently been the subject of much
debate and experimental work.  It was previously believed that the regen-
eration of the OH radical proceeded via the reaction

                               H + H02 •* 2 OH                         (4)

Davis et al.  (1973) have concluded that reaction 3 proceeds quite slowly
                                   3  3
(k  probably being less taan 6 x 10  m /mol'sec) and therefore, should not
play a significant role in atmospheric chemistry.  Westenberg though
believes that reaction 3 probably proceeds even faster than reaction 1.
Westenberg and de Haas (1972) derived a rate constant for reaction 3 at
room temperature of about 6.02 x 10  m /mol*sec which is indeed faster
than that of reaction 1.  Explanations offered to explain the difference
in these k  values (Davis, Payne and Stief, 1973) are not conclusive.
          O
It has been pointed out  (K. L. Demerjian, 1974) that the rate constants
for reaction 3 are erroneous.  For further discussion of these reactions
and many more of interest the reader is referred to a recent work by
Demerjian (1973).
     Naturally, if this mechanism is correct it would require that a
significant amount of hydroxyl radicals be present.  Assuming that CO
is in a steady-state in the atmosphere with a concentration of 100 yg/m
(.1 ppm) Weinstock (1969) calculated the OH required to maintain this
steady state condition according to the equation:

                              [OH] = P/k[CO]

where P = production rate of CO and k is the second order rate constant.
                             -3         3
He calculated that 1.16 x 10   mol OH/m  were required and that requirement
could easily be satisfied by an OH regenerative process as suggested above.
-------
                                    -59-
     Vertical CO profiles of the atmosphere have recently been measured

by Seller and Junge (1969), Seller and Warneck (1972)  and Goldman,  Murcray,

Murcray, Williams, Brooks and Bradford (1973) .  These  first two reports

both found a sharp decrease in the CO mixing ratio when crossing the

tropopause into the stratosphere.  The mixing ratio is defined as the

mass of gas per unit mass of dry air.  Goldman et al's. profile did not

show a sharp decrease in the mixing ratio at the tropopause but they

believe this might be due to local atmospheric conditions just around

Holloman, New Mexico,  where their measurements were taken.  It seems

obvious from vertical  profiles such as these that there is a significant

decrease in CO levels  above the tropopause, thus supporting the conclusion

that the stratosphere is a sink for carbon monoxide.

     Quantitative estimates of the CO destroyed in the stratosphere by

this reaction mechanism have not been made, although Pressman and

Warneck (1970) believe that virtually all CO entering  the stratosphere

is so destroyed.  Therefore, the size of the stratospheric sink is

dependent upon the transport rate of CO rich air through the tropopause

into the stratosphere.  They have estimated this flux  at 1.3 x 10
     2
mol/m *sec but admit the degree of uncertainty in this estimate is very
high due to insufficient data.  Based on this flux and recent estimates

of the total CO reservoir in the troposphere,  they estimated that about

11-15% of the total CO inventory in the troposphere is destroyed in the
                                                                          9
stratosphere by the reaction sequence mentioned above, or about 52-71 x 10

kg/yr.

     Although the above reaction sequence has been recognized as a major

sink in the stratosphere, only recently has it been suggested to be a
-------
                                    -60-
sigrtif leant destructive reaction in the troposphere  (Weinstock,  1969) .



Recent investigations (Levy, 1971; McConnell, McElroy  and Wofsy,  1971)



have shown that sufficient amounts of hydroxyl and hydroperoxyl  radicals



exist in the troposphere to provide a mechanism for  CO oxidation there



as well .



     Levy (1971) suggested the following reaction sequence  which provides



for both the production of hydroxyl and hydroperoxyl radicals  and the



oxidation of CO and NO in the troposphere.
     (1)  0  + hv  (2900A • 0('5P) + M


                                      73    -1     -1
                        k  = 3.01 x 10  m  mol    sec
      (3)  OD) + HO -> :>OH



                        k  =  1.8  x  10  m   mol    sec
      (4)  OH + 03 -> H02
                                       53     -1     -1
                         k  £ 3.01  x  10  m   mol    sec
      (5)  OH + CO •+ C02 + H
                                      43     -1     -1
                          p =  9.0  x 10  m  mol   sec
      (6)  H + 0  + M ->- HO  +  M
                         k,  =  1.08  x 104 m6 mol"2 sec 1
                          D




      (7)  HO  + NO  -*• NO  f  OH
            ^           z


                         k  =  3.01  x 10  m  mol    sec





More  recent estimates  of these rate constants by Kummler and Bauer (1973)



are in  excellent  agreement  with those of Levy listed above.
-------
                                     -61-
Although quantitative estimates of the total CO possibly consumed by




this reaction series have not been made, Levy did calculate a carbon




monoxide lifetime in the surface atmosphere of 0.2 years using hydroxyl




radical concentrations which he had determined.  This estimate is in good




agreement with that made by Dimitriades and Whisman (1971), but is still




the subject of considerable dispute among other scientists (Kummler,




Bortner and Jaffe, 1971) who produce their own evidence to show that gas




phase kinetics at ground level cannot alone explain the sink anomaly.






Carbon Monoxide Cycle




     The last several years have seen an explosion of interest in




defining natural CO sources and sinks.  While several researchers have




confirmed that soil and the reaction of CO with hydroxyl and hydroperoxyl




radicals in the stratosphere and possibly the troposphere constitute




a major sink for CO, other researchers have found additional natural




sources of this gas.  These sources, which include oxidation of methane,




oxidation of terpenes and the oceans to name just a few, are now thought




to contribute more CO to the atmosphere than that emitted from anthropo-




genic activity.  This is an outright contradiction of what only two years




ago was thought to be the final word; i.e., anthropogenic CO emissions




are many times greater than natural CO emissions.




     Table IX compares recent estimates of the importance of CO sources




and sinks.  It is obvious from this table, if these estimates are at all




reasonable, that soil and gas-phase oxidation in the stratosphere and




troposphere might be capable of consuming all the CO that nature and man




produce.  If so, this explains why the background concentration of CO




has not increased over the last several decades, and we might suggest



that there no longer exists a "CO sink anomaly".
-------
                                   -62-
    Table IX.   ATMOSPHERIC FLUXES OF CARBON MONOXIDE (in 10  kg/yr)
I.   Sources

    a.   anthropogenic

    b.   natural

        1.   oceans
                                      359
                                   43-220
        2.  oxidation of terpenes     54

        3.  oxidation of methane     900

        4.  other              	1_

    TOTAL                      1356-1533
Jaffe (1973)



Liss and Slater  (1974)-
Linnenbom et al.(1973)

Robinson and Moser  (1971)

McConnell et al.  (1971)
II.   Sinks

     a.   soil                     67-1400

     b.   gas-phase oxidation

         1.  stratosphere           52-71

         2.  troposphere        	?_

     TOTAL                       119-1471
                                             Heichel  (1973)-
                                             Ingersoll  (1972)
                                             Pressman and Warneck  (1970)
-------
                                    -63-
                         NITROGEN CONTAINING GASES





     Although the most abundant oxide of nitrogen in the lower atmosphen



is nitrous oxide (N«0),  it does not play an important part in air pollu



tion chemistry.  The two nitrogen oxides which are important are nitric



oxide (NO) and nitrogen dioxide (NO ) for they play a prominent role in



the generation of photochemical smog.  In fact, the presence of nitrogen



dioxide, as a light energy acceptor, is a necessary condition for



photochemical reactions to proceed.



     Measurements of the background levels of nitrous oxide consistently



average to about 460-490 yg/m  (0.25-0.27 ppm) (Schutz, Junge, Breck



and Albrecht, 1970).  Ambient N09 levels are much lower, probably around


      3                                                    3
2 yg/m  (.001 ppm) (Robinson and Robbins, 1968) to 2.6 yg/m  (.014 ppm)



(Junge, 1963).  Too few accurate measurements have been made of ambient



levels to obtain a meaningful average, but Lodge and Pate (1966) found



concentrations up to 11 yg/m  (.006 ppm) in Panama.  Not many measure-



ments of NO levels have been made but they are approximately the same



as those of NO .



     The other nitrogen compound considered is ammonia.  Although ammon ;.<"



per se, is a relatively unimportant air pollutant, it does play an



important role in atmospheric chemistry through its part in the formation



of aerosols.  Ambient concentrations of ammonia seem to average about



4 yg/m  (.006 ppm) (Robinson and Robbins, 1968).
-------
                                    -64-
NTTROGEN OXIDES





Sources



     Nitrous oxide (N-0.),  the most abundant nitrogen compound in the



atmosphere, is produced from bacterial decomposition of other nitrogen



compounds within the soil.  Arnold (1954) was the first to study the



production of N00 by soils.  He found that the factor most conducive
               t~


to increased NO evolution is a high soil moisture content, especially



if a nitrogen source, stch as nitrate or ammonia, is present.  Bacterial



decay is the primary source of nitrous oxide and so is probably responsible



for ambient N?0 concentrations.  An estimate of the flux of N90 from
             <->                                               ^


soils into the atmosphere has recently been made by McConnell (1973).



He estimates that although an unknown amount is released and reabsorbed,



about 1.1 x 1010 kg N 0 - N (1.73 x 1010 kg N20) are available for oxida-



tion or photolysis to NO  and N  in the stratosphere each year.  Schutz
                        .\      L»


et al.  (1970) declined to extrapolate the data they collected on three



soil samples to attain a global flux.  Their measurements showed a flux

                   Q        7

on the order of 10   g N/;)0/m  'sec, an order of magnitude which, if



maintained globally, would necessitate an NO residence time of about



70 years.  This is the same residence time they estimated based on the



photodissociation rates of N~0 in the troposphere and stratosphere.



Goody and Walshaw (1953) estimated a globalN?0  production  rate  of  about


        12
100 x 10   kg/year and Robinson and Robbins  (1968) suggested that  soils


                       10
produce about 59.2 x 10   kg  NO each year by biological action; and



of this about 55.4 x 10   kg  (35.3 x 10   kg NO - N) are reabsorbed



by the  soil and about 3.8 x 1010 kg N20  (2.4 x  1010 kg N20 - N) travels



up to the stratosphere where  it is destroyed.   It is this  latter rate



that  is shown in Table XI.
-------
                                    -65-
     Craig and Gordon (1963) raised the possibility that the ocean



might be a source or a sink of NLO.  Bates and Hays (1967) concluded



though that the uptake of N^O by the oceans in areas where upwelling



waters are deficient; in NO is of negligible importance.  Based upon the



difference between the mean N_0 concentrations of oceanic surface waters



and adjacent air, and an estimate of the total liquid phase resistance,



Liss and Slater (1974) estimated the flux of N?0 from the ocean to the



atmosphere at 1.2 x 10   kg/year.  This flux is based entirely on theory,



but indicates that the oceans could conceivably release significant



quantities of N?0 to the atmosphere.  Laboratory and field experiments



are needed to decide whether or not the ocean is a source of N?0.



     Production rates of NO and NO  by soils are much more difficult to



measure or estimate, and good data are lacking.  McConnell (1973)



recently summarized a few of the problems involved in appraising the



amount of nitrogen oxides produced by soil.  He contends that this soil



source is small compared to that produced as a result of the gas phase



oxidation of atmospheric ammonia by OH which he thinks produces 7 x 10



kg NOY - N per year.  McConnell offers alternative reaction sequences
     A


for NH  in the atmosphere; one reaction sequence provides a constant



source of NO, the other a sink.  If the latter is shown to occur in



the atmosphere, in order to account for the amount of NO known to be in



the atmosphere, then an additional source of NO must be found.  In this



case, McConnell concedes that the soil might actually constitute a



significant source of NO , on the order of 10   kg/yr.
                        A.


     Nitric oxide is also produced as a result of the photolysis of



N_0 in the stratosphere as follows:
-------
                                    -66-
                             N'20 + hv -»• NO + N




                                                                       -9
The photolysis rate for this reaction is less than or equal to 7.4 x 10



sec   (McElroy and McConaell, 1971).  Whereas these authors estimate the

                          o

production of about 3 x 10  kg NOY - N per year in this manner, Bates
                                 A


and Hays (1967) estimated that 3.5 x 10   kg NO are produced annually



by this photolysis reaction.



     The other primary source of nitrogen oxides is anthropogenic,



primarily combustion processes.  Estimates of production rates for NO



and N0? are included together because available emission data rarely



distinguish between these two forms.  Robinson and Robbins (1970)


                      "                      9
estimated that 53 x 10  kg of NO , or 16 x 10  kg NO- - N were emitted



annually (here again NO  includes both NO and N0? production).



     I_n toto, natural emissions of NO   (including NO) are approximately


                                                       9               9
15 times greater than anthropogenic emissions (768 x 10  kg vs. 53 x 10



kg N00) (Robinson and Robbins,, 1970).  Therefore, anthropogenic emissions



play only a minor part in the total circulation of nitrogen compounds



in the atmosphere.





Removal Mechanisms



     At 20°C and 1 atm total pressure, the solubility of NO in water



is 0.121 g/lOOg HO.  Upon release from soils, an unknown portion is



believed to be removed by vegetation, soil and water.  No information  is



available on these reactions.  The major portion of the released N?0



is destroyed by photodissociation in the stratosphere and upper tropo-



sphere.  Under normal tropospheric conditions, N-0 is chemically inert



and it partakes in no other  chemical reactions.  The residence time of
-------
                                     -67-






nitrous oxide is probably around  70  years  if the,r-



biosphere (Robinson and Robbins,  1968), but  coul'i



years if there is a biologic  loss mechanism.   Hid,  i i'i



residence time as 4 years.



     Nitric oxide is rather insoluble  in water.      >!:



pressure, its solubility is 0.00618  g/lOOg HO   •'



the other hand, immediately dissociates in walti  •



For this reason the solubility  of N0_  in watei i -   • :•



     Nitric oxide is either oxidized to NO  or [/



is then removed primarily be  precipitation,  more  ••'  • ••



form of nitric acid (HNO ).   It can  also be  abbu^  ;



soils or participate in photochemical  reactions  :    »•



 their reactivity, the residence timesof NO  and  '••' .••••



probably around 5 days (Hidy, 1973) .   Nitrogen *' .  >•• "



the atmosphere by the following mechanisms.





     Vegetation:  Vegetation  has been  shown  cap.;!"''  >>!



cant amounts of NO  and NO from the  atmosphere,   i; •.<



that alfalfa and oats absorbed  NO  from the  aij


     2
mol/m -sec when exposed to an atmosphere contain •>•••  !;-


       -5      3
M x 10   mol/m ).  With time,  the uptake  rate u-   <*•



to air containing between 300 and 460  yg NO,,/in  '••  "  •



uptake rate remained constant when the plants v/u-'  •  •



150 ug NCL/m .  He also found that there was  nn  >'( < • •  •



rates when the intensity of the sunlight was  ivi •  "';.



     Tingey extrapolated his  data to estimate !L-  ;-.•""•



the Salt Lake Valley in Utah.   His estimates, iv-.-;  -«
-------
                                    -68-
                2
of about 2000 kin  in the valley with an 85% vegetal covering in the summer,



are tabulated below (Table X).
            Table X.  WTES OF NO  UPTAKE IN SALT LAKE VALLEY, UTAH
                                Removal Rate               „     , _.
              ,,                          „                 Removal Rate
          NO^ Cone.                    ,2                  ^    ,, 11
            2                      per km                   from Valley

      (24-hr average)             (kg/day)                    (kg/day)
          40 yg/m3                  18                      36 x 103



          20 yg/m3                   9                      18 x 10



          10 yg/m3                   4.5                     9 x 103
     On an annual basis, removal of NO  from the valley with an ambient



NO  concentration of 9.6 yg/m   (about 3f-4 times greater than background



NO  levels) would total about 3 x 10  kg N0?.  This may be compared with



the estimated total global anthropogenic annual N0? emissions of



53 x 109 kg (Robinson and Robbins, 1970).



     Hill  (1971) found in his experiments on the uptake rate of gases  by



an alfalfa canopy that NO was absorbed with  a depositional velocity of



0.1 cm/sec.  NO  was absorbed at a velocity  of 2 cm/sec when present



in the air of the chamber at a  concentration of 2  x 10   mol/m   (or



96 yg/m ).  Using ambient N02 concentrations found in thbse areas  of



Southern California from August to October of 1968, and assuming a



continuous alfalfa cover, Hill  estimated that N02  could be removed at

                 T

a rate of  0.1 g/m^'day.  Because nitrogen is often a  limiting factor  in
-------
                                    -69-
plant growth, Hill believes his study  shows  that  NO,  and  ;j'i rc



as much as .03 g N/m -day.  Since NO   dissociates in  walei   .i



that the N absorbed from the atmosphere  as NO  would  be me I •''••



so the amount absorbed over one year could be  significant
     Soil:  Nitrogen oxides  (especially  N_0)  have  long  be- ;  I
        "'""                                 Z*


produced by biological action in soils.   Recently  though,



(1971) found that soils could absorb nitrogen dioxide fivn



as well.  They found that when they passed  air containing



in a test chamber, the concentration of  NO   in the air  wa.,  .


                             3     3
an initial value of 190 x 10  yg/m   (100 ppm)  down to S, '    !



(3 ppm) over a 24 hour period.  When the soil  was  autoci;



NO-present over the same time period was reduced from J8(    '



(97 ppm) to only 25 x 10  yg/m   (13 ppm) .   The reaction l.\ '• <



be the cause of the NO- uptake in the  soil  was not discovt



Extrapolating the results of this experiment,  the  author.:,    •-•,



the soils of the United States might be  capable of remcn i; -



of NO- per year from the atmosphere, an  amount they point ••>•''-



bit under 20 times the total annual production of  NO, 311  i i-: •



(3.3 x 10   kg).  Of course, such a removal rate is larg'-i



for there are too many variables that  need  to be taken  ui: •   .



such as soil moisture content, soil composition, both iiuix,",.:1



organic, and vegetal covering to name  just  a few.   AithouHi  T



such an estimate is warranted, the limitations placed upon  > 'i



must be kept in mind until more information becomes avajJ, :M



     Nelson and Bremner (1970) point out that  the  NO ahst .  -



will ultimately be oxidized to nitrate.   These nitrates c <  <
-------
                                    -70-





decompose and result in the production of nitrogen dioxide again.  The


                                                              -4
rate of N0? production by nitrate breakdown in soils is 2 x 10   g NO /


 9

m"*hr (Marchesani, Towers, and Wohlers, 1970, and Makarov, 1970, as cited



in Bohn, 1972).  This NO  production rate is dependent upon the nitrate



content of the soil and does not proceed during darkness.



     Nitric oxide may also be absorbed by soils, but is then oxidized



almost immediately to NO  (Mortland, 1965; and Bremner and Nelson,



1968).  Mortland has also discovered that transition metal ions in



the soil promote NO absorption.  If the soil is saturated with alkaline



earth cations though, absorption of NO is halted.  Sundareson, Harding,



May, and Henrickson (1967) found that alkaline-earth zeolites readily



absorb NO and release it as NOY and HNO_ when heated.  To date, the
                              A        o


role organic matter plays in the absorption of nitrogen oxides by soil



remains a mystery.  Ganz, Kuznetsov, Shlifer, and Leiken  (1968, as cited



in Bohn) found that upon passing NOY-contaminated air through 1 meter of
                                   A


peat,all nitrogen components were removed.  Organic matter is such an



important component of soil that to not be fully aware of its effects



on a gas could only hinder understanding of the mechanism of absorption



by the soil.  Obviously, nore research is needed in this area.





     Water Bodies:  There are very little data available on the amount of



nitrogen oxides absorbed by the oceans.  Craig and Gordon  (1963) first



suggested that the oceans might constitute a sink for N_0 when they found



that the sea water at depth was depleted in N?0 compared to the surface



waters which they found tc be in equilibrium with atmospheric concentra-



tions.  This depletion could not be explained by temperature variations



with depth.  Based on the mean upwelling speed of the ocean's waters,
-------
                                    -71-




                                         -4      -5
thought by Bowden (1965) to range from 10   to 10   cm/sec, Bates and



Hayes (1967) estimated the potential sink strength of the oceans for N?0


                          -7              6      ?
as ranging between  6  x  10   and  6  x  10~  kg/m -yr-   •   Ljss and Slater



(1974), on the other hand, concluded that the flux of N90 uas from the sea



to the air.  They calculated a flux rate of 3.2 x 10  kg/m""yr based upon



an N«0 concentration gradient across the air-sea interface measured by



Junge and Hahn (1971)  .  The total flux of NO from the sea to the



atmosphere, if assumed constant over the entire oceanic surface, is



1.2 x 10   kg/yr.  With such an obvious contradiction as to the direction



of the NO flux,  there is a real need for additional research in this area.





     Washout and Rainout:  The major sink identified thus far for nitrogen



oxides (NOY) is the solution of soluble species in cloud anil rain droplets
          A


with subsequent removal by precipitation.  According to Haagen-Smit and



Wayne (1968) the reaction is:
                         4NO
Georgii (1963) believes the reaction proceeds as follows:
                        2NO  + HO -v HNO  + HNO
                           £•    £.       O      .£.



                                         3

Georgii found when a total volume of 20 m  of atmospheric air containing



average concentrations of NO , but cleansed of all natural aerosol



particles, was  passed over both distilled and rain water at a velocity



of 5-10 cm/sec, that the nitrate (NO ) concentration in the rain water
                                    O


increased from 0.18 to 0.45 mg/1 while the nitrite (N0~) concentration



increased from 0.04 to 0.08 mg/1.  Georgii further showed that the N0~
                                                                     3


concentration in rain is related to the concentration of NU  in
-------
                                   -72-
the atmosphere.  In Frankfurt am Main, Germany, where ambient NO  levels



are higher in winter than summer, Georgii found that the nitrate content



in rainwater was also greater in winter than summer.



     Robinson and Robbins (1968) suggested that the dissolution of NO



proceeds as follows:






                          3NO,, + HO + 2HNO  + NO
                             £*    £.        J





However the hydrolysis reaction proceeds, the outcome is the same in



all cases; the nitric acid formed is absorbed onto hygroscopic particles



or reacts with atmospheric ammonia to form nitrate salt aerosols (NH NO



for instance).  It is thea either removed by precipitation, or if



vaporization of the drool3t occurs, by dry deposition.



     McConnell  (1973) estimates that 2 x 10   kg NO" - N is removed from



the atmosphere each year by precipitation, and an additional 7 x 10   kg



NOY - N removed by dry deposition, the major portion of this probably
  A


being HNO  .





     Atmospheric Reactions Involving Nitrogen Oxides:  Although more



nitrous oxide  is released to the atmosphere than any other nitrogen



oxide it does not play a major or very complex role in atmospheric



reactions.  Because it is chemically inert in the troposphere, its sink



lies in the stratosphere where it is transported by vertical mixing and



destroyed  by photolyzing reactions.  Bates and Hays  (1967) indicate



that the most  significant reactions are:






                     N20 -- hv ->- N2 + 0  (1D)  ; A < 3370 A






                     N20 i- hv •> NO + N  (4S)  ; A < 2500 A
-------
                                    -73-
The latter reaction they consider responsible for about 20% of the total



dissociation in the stratosphere.  They calculate an average latitudinal


                                 -9    -1
dissociation rate of about 3 x 10   sec



     Nitric oxide can be removed from the atmosphere by several reactions.



The primary reaction is its oxidation by ozone to form NO
                              NO + 03 + N02 + 02




                                      18                ^     1
which has a rate constant of 9.0 x 10"   exp (-1200/T) m  sec"   (Schofield,



1967).  In the upper atmosphere, NO can be photolyzed to atomic N which



may react with other NO molecules to form N  as follows:





 NO + hv -> N ( S) + 0 ( P) ; photolysis rate = 4.0 x 10   sec    (Callear

                                                      and Pilling, 1970)





  N + NO + N. + 0;k = 2.2x 10~U cm3 sec"1 (Schofield, 1967)
     Nitrogen dioxide also can engage in a number of reactions.  It may



be oxidized to gaseous NO, by the reaction
                         «5





                              N02 + 03 + N03 + 02




                                  •JO                -7

The rate constant equals 9.0 x 10"   exp (-3500/T) m /sec  (Schofield, 1967).



Or, it may form nitric acid by its reaction with hydroxyl radicals as



follows:






                             OH + N02  j HON02* -> HN03






The nitric acid formed is removed by precipitation.   For further elaboration,




the reader is referred to McConnell and McElroy's (1973) article.




     The importance of NO and NO  as pollutants reflect their participation



in photochemical reactions.  In polluted atmospheres they react with SO-
-------
                                    -74-
and hydrocarbons to forn aerosols.  Probably the most important photo-



chemical reaction involving NCL is its photodissociation as follows:





                    N02 + hv (2900 A < A < 3800 A)  t  NO + 0





This atomic oxygen then is free to react with molecular oxygen and  form



ozone.



     Peroxyl radicals  (ROO)> formed between reactive free radicals  (R)



and CL, can react with NO and NO  to form alkyl nitrates or peroxyacyl



nitrates as follows:





                        ROO* + NO ->• ROONO- h^ R0« + NO





                        ROO' + NO  -»• ROONO
                                 ^-        L*




These secondary reaction products are targets of further photochemical



attack.  For example, peroxyacyl nitrate may be photodecomposed into  NO



and acylate radical as follows:




                          0           0

                          11       hv  "
                         RC-OONO  +  RCO- t N02





This rather brief overview of possible photochemical reactions involving



nitrogen oxides is abridged from Haagen-Smit and Wayne  (1968) .  It  is



inappropriate here to  elaborate on photochemical reactions involving



nitrogen oxides.  The reader is referred to Altshuller and Bufalini (1971),



Cadle and Allen  (1970), and Leighton  (1961) for more detail.





Environmental NO  Cycle
                A


     The circulation of nitrogen oxides in the atmosphere is  complex



and not well understood.  The global NOY cycles formulated by Robinson
-------
                                    -75-

and Robbins (1968, 1970) and McConnell  (1973) make clear how very little
good quantitative data exists.  A more complete global nitrogen cycle
(Figure 4) for both nitrogen oxides and ammonia is presented at the end
of the following section on ammonia.
     The holes in these cycles are obvious and major ones (see Table XI).
We need, for instance, a better idea of how much NO and NO,, is released
from the soil, how much NH  is oxidized to NOY, and how much NO and N00
                          J                  A                        /
is destroyed by photochemical reactions.  We also need to know why large
discrepancies exist between estimates, and the bases for these estimates.
For example, McConnell bases his estimate of NO  removed by dry deposi-
                                               A
tion on soil data, the nitrogen content of precipitation and a deposition
velocity factor.  Robinson and Robbins, on the other hand, use the
deposition velocity function to estimate gaseous deposition, a removal
mechanism McConnell makes no mention of.  These and other inconsistencies
indicate the need for further study.
-------
                                    -76-
                                                                10
Table XI,   ATMOSPHERIC FLUXES INVOLVED IN VARIOUS NOY CYCLES (10   kg N/yr)
                                                    A

I, Sources
a. anthropogenic; NO, NO
b. biological; N?0
c. biological; NO, N0?
d. oxidation of NH
e. stratospheric transport;
NO, N02
TOTAL
11. Sinks
a. rainout
b. dry deposition
c. oxidation of N20-»NO (strat.)
d. photolysis of N20->N2 (strat .)
e. gaseous deposition
TOTAL
Robinson and
Robbins
(1968)

1.5
1.2
30.4
N.E.
	

33.1

112.9
27.1
0.2
	
10.7
150.9
Robinson and
Robbins
(1970)

1.6
1.2
23.4
N.E.
	

26.2

7.5
1.9
0.2
	
4.5
14.1
McConnell
(1973)

1.8
1.1
?
7
0.07

9.97

2
7
0.03
(1.07)
	
(10.1)
N.E. = Mechanism recognized but no estimate made.

*NOTE:  Friend  (1973) has pointed out that due to an error in converting units
        of kilograms per hectare to tons per square meter, much of Robinson and
        Robbing nitrogen compound cycle is invalid.  Friend though, gives no
        indication as to which values are wrong.  Due to lack of information,
        their cycle is nonetheless shown for comparison.
-------
                                    -77-
AMMONIA





Sources



     Most ammonia in the atmosphere is the result of the bacterial



decomposition of organic material on the earth's surface.  The factors



which affect the emission of this NH  from the soil are its nitrogen
                                    O


content, pH, and moisture content (McConnell, 1973; and Georgii, 1963).



NH  is more readily released from dry soils than moist ones, and is more



readily released when the pH of the soil is greater than or equal to 6.



It has been suggested that in acid soils NH. is probably biologically



oxidized (autotrophic nitrification) to NO  and therefore not  available
                                          ij


for release.  The net oxidation reaction would be





                        NH* + 202+  NQ~ + H20 + 2H+





(Keeney, 1973).  DuPlessis and Kroontje (1964), on the other hand,



suggested that the low hydroxyl ion activity in acid soils and adsorption



of NH  on soil colloids prevents or retards the release of ML, according



to the reaction
                        NH* -H OH~^r  NH3 -*- HO




In alkaline soils the activity of OH  is greater than that of H  and so



NH  is readily released.



     Junge (1963) suggested that the oceans may contribute some NH  to



the atmosphere, but to date there has been no accurate measurement made



upon which to base an estimate of the source strength.  McConnell has



estimated that these biological sources release about 17 x 10   kg NH  -



N each year,  whereas Robinson and Robbins ' (1968) nitrogen cycle calls
-------
                                    -78-




                                         12
for the release of NH, on the order of 10   kg NH_ each year (see



Table XV),



     Anthropogenic NH  emissions result primarily from the combustion of
                     O


coal.  Robinson and Robbins (1970) estimated that the atmospheric ammonia



burden due to man's activities is 0.4 x 10   kg NH» - N per year, less
                                                  «D


than 2 1/2% of the estimated NH  burden due to biological emissions



(17 x 1010 kg).
Removal Mechanisms



     Ammonia is extremely soluble in water.  At 20°C and 1 atm its



solubility in water is 62.9 g/lOOg H?0.  Ammonia therefore can be



readily absorbed by water bodies and vegetation.  It is also removed



by its reaction with hydroxyl radicals to form nitric oxide.  Probably



the most important removal mechanism though is its solution in rain



water along with S0~ or other gases to form aerosols such as (NH,)7SO..



The quantity of NH  removed by precipitation is ordinarily determined by
                  *J


measuring the quantity of NH. present in rain.  From the data available,



the residence time of ammonia in the atmosphere is probably around 7 days.



The following mechanisms are important in removing ammonia from the



atmosphere.





     Vegetation:  Because it is normally present in the atmosphere in



such small concentrations, ammonia's interaction with soil, water and



plants has long been ignored.  Hutchinson, Millington and Peters  (1972)



and Porter, Viets and Hutchinson  (1972) have found that even at the low



naturally occurring atmospheric concentrations, plant leaves absorb



significant quantities of ammonia.  Porter et al . exposed air containing
-------
                                    -79-
  N labeled NH  to  corn seedlings in growth chambers.  '!'!•<'  -<.<\i  root
              O


system was isolated from the atmosphere by polyethylene bup  i?'i>ch were



tested and found to be  impermeable to NH .   At the condition of" the testing
                                         O


period, the plants  were analyzed and found to contain  "N  ,v-i "."od nitrogen



compounds.  They therefore deduced that direct foliar absorp1 i--n was



responsible for the presence of most of the   N found in f.h--  :>' nt.



They have indicated that simple isotopic exchange might a<:,c»;iut  tor a



small part of the apparent [  N] NH  uptake.  They did not  att'.-mpt to



establish an uptake rate for NH  by these plants.



     Hutchinson et  al.  (1972)  based their estimates of Id1,'  u|.t >K; rates



of various plants on the difference in the mass flows oF iv|i   ,u-'. ,i-..js-ed



absorption rates of NH   showed large diurnal fluctuation- .   liii/-  « '••[< --uecies



and for soybeans at three different nitrogen fertility lev."





     Soil:  The absorption of NH  by soils has received rvi.'i rv.'y little



attention until recent  years.   Malo and Purvis (1964) ijiV".i ii;i'fd the



seasonal variations in  atmospheric NH  concentrations over  K'.1^ Jersey in
                                      O


relation to the rate of NH  absorption by six New Jersey '.^"is.   The
                           •J


atmospheric NH  content averaged 57 yg/m ,  considerably -i>(.•••[<>-<  i ban the



normal background concentration of about 4.2 yg/m .   The fM. v..-.  dtmospherJc
-------
                                    -80-
Table XII.   AMMONIA ABSORPTION RATES FOR FOUR CROP SPECIES AND FOR SOYBEAN

            AT THREE NITROGEN FERTILITY LEVELS
Added
„ . soil
Species
^ nitrogen

Soybean (Glycine max)
Soybean (Glycine max)
Soybean (Glycine max)
Sunflower (Helianthus annuus.)
Corn (Zea mays)
Cotton (Gossypium hir^utum)
(mg)
0
5
20
l 5
5
5
Leaf
area
(cm )
65
80
84
96
58
55
NH3 in
chamber air
(yg/m3)
29
24
24
31
24
44
NH3
uptake
rate,
(yg/m -hr)
410.
420.
400.
490.
560.
350.
NH_ in New Jersey is probably due to the combustion of coal, oil and
  J


gasoline.  The rates of absorption by six soils under field conditions



in New Brunswick are given in Table XIII, along with some soil character-



istics.



     The factors governing the amount of NH  absorbed by soils were
                                           O


investigated by Hanawalt  (1969a,b).  Hanawalt  (1969a) tested the



influence of several atmospheric factors on the sorption rate which



included the length of exposure, the ambient NH, concentration, the



ambient air temperature and the velocity of the air passing over the



soil.  The most important factor causing a change in sorption rates was



the ambient NH  concentration.  As the concentration increased from
              O
                     3                         3
about  37 yg NH  - N/m  to about 67 yg NH  - N/m , the sorption rates
              J                         J
-------
 Table XIII,
                                     -81-
SOIL CHARACTER AND ABSORPTION RAThS
FIELD CONDITIONS
Date and Soil
November 21, 1962
1. Collington sandy loam
2, Dutchess loam
3. Lakewood sand
February 18, 1963
4. Matapeake silt Ica/i
5. Dutchess shale loam*
6. Nixon loam*
PH

4.6
5.4
4.0

5.6
7.0
5.7
Organic
matter

2.32
7.12
2.51

2.92
--

q
0

li
26
0

Ifa
-
_
                                                     I
                                                     I




                                                 O.t. :
 *Soils 5 and 6 both from cultivated fields

**Cation Exchange Capacity (C.E.C.) in milliequivalu-




 of the six soils tested were found to increase by ti'  •


 also found increasing sorption rates to be correiai  <


 temperature, wind velocity or exposure time.


      Hanawalt (1969b) evaluated the influence of st


 such as organic matter, pH,  cation exchange capacit •-•


 moisture on NH  sorption rates.  The property found •


 sorption rates was the soil  moisture content, betwi.••••'


 a positive correlation.  Hanawalt concluded then in•*
 Nil,  is  absorbed by a soil is affected most "by thu
-------
                                    -82-






the rate of supply of anmonia to the soil surface," namely the atmospheric




factors mentioned above, and "the fraction of the ammonia molecules which




are actually captured," which is dependent upon the moisture content




and air permeability of the soil.






     Water Bodies:  Because the absorption of atmospheric ammonia by lakes




and streams can promote eutrophication in that water body, a great deal of




concern has developed, especially in areas adjacent to large cattle




feedlots and sewage treatment plants which both release great quantities




of ammonia into the atmosphere.  Hutchinson and Viets (1969) showed that




the ammonia volatilized from cattle feedlots was absorbed by nearby




water bodies.  In their study, they calculated that Seeley Lake in




northeastern Colorado, which is about 2 km from a 90,000 unit feedlot,




absorbs enough Nil  from the air over a one year period to elevate its




nitrogen concentration by 0.6 mg/liter.  An interesting result of their




study was the discovery that although there was a tendency for the NH




content in precipitation to increase as the feedlots were approached,




in no instance at any site, near or far from a feedlot, was the NH




content in precipitation very high.  It appears that in all cases the




amount of NH  removed from the atmosphere by precipitation was "insignifi-




cant compared to the amount absorbed directly from the air by aqueous




surfaces in the vicinity of cattle feedlots"  (see Table XIV).




     In an attempt to quantify the removal of NH  released from a point




source  (such as from a  plume produced by a sewage treatment plant) by  an




aqueous surface, Calder (1972) developed a simple mathematical model




which takes into consideration the atmospheric transport and diffusion




o£ the plume as well as a characterization of the removal process,
-------
                                   -83-
Table XIV.  NH* - N ABSORBED BY SURFACE WATERS AND ITU
            ADJACENT TO A CATTLE FEEDLOT
* Mean weekly absorption of ammonia-nitrogen by di!
  the period 27 July 1968 through 27 February 1969
  where measurement began 27 September 1968)
t Estimated annual absorption of NH  - N by water ^
  dividing mean weekly rates by two and multiplying
  are divided by 2 because that is the ratio of NH^
  acid traps to lake water traps)
§ NFL - N in precipitation measured during the perj-.
  21 November 1968,  Total amount of p'pt. over thj:
  site 1 and between 3,0 and 3.6 cm at the other sii
  of NH* measured in the precipitation was extrapoJ,
  to give the measurements listed.
Site                   Site description
  1  No feedlots or irrigated fields within          i> '        "      •'  .:
     3 km; no large feedlots or cities within 15 km

  2  Only small (less than 200-unit) feedlots
     within 4 km,  none closer than 0.8 km

  3  About 0.2 km east of 800-unit feedlot and
     0,6 km west northwest of another similar
     feedlot

  4  On northeast  shore of Clark Lake and 0.5 km
     southwest of 9,000-uait feedlot

  5  On southeast  shore of Seeley Lake and 2 km
     west northwest of 90,000-unit feedlot

  6  About 2 km east of 90,000-unit feedlot          ;        '•           to

  7  About 0,4 km west of 90,000-unit feedlot          '                 -I
-------
                                   -84-
in this case the deposition velocity.  The reader is referred to his


article for the mathematical development of this model.  Given the



uncertainty in the value of the deposition velocity for NH, absorption


by water, Gaidar concludes that absorption of NH  from an NH, cloud,
                                                »J           *5

passing over a water body some 30 km long at an average speed of 5 m/sec,


could exceed 20 per cent.




     Rainout and Washout:  Ammonia, as previously mentioned, is extremely



soluble in water.  Once dissolved it ionizes to NH., the amount depending


on the pH, as follows:




                            .  NH3'-H20 £ NH* + OH"




The rate constant equals J..774 x 10   sec   (Robinson and Stokes, 1959).



Normally the pH of cloud water ranges from 5 to 6.  Under atmospheric


conditions, assuming that an equilibrium solution is attained and CO  is


present in normal concentrations, Junge  (1963) has shown that approxi-


mately 87% of the NH, present in air at a concentration of 3 yg/m  could


be absorbed in cloud droplets.  Unfortunately, equilibrium is not usually,


if ever, reached in cloud droplets.


     Ammonia has been shown to be an important catalyst for the oxidation


of SO  and NO  in solution  (van den Heuval and Mason,  1963).  Both Scott


and Hobbs  (1967] and Miller and de Pena  (1972) have shown that as the



partial pressure of NH   in the atmosphere increases, greater concentra-



tions of SO  can be dissolved and oxidized in solution to form sulfate



particles.  The resulting aerosols,  if evaporated, are composed largely


of  (NH ) SO  particles with probably minor amounts of  NH HSO,, NH HSO.


and  (NH4)2SO   (Miller and de Pena, 1972).  In the case where NH  is
-------
co-absorbed with NO , the resulting particles would be composed



primarily ol NH NO .



     It has been estimated (Robinson and Robbins, 1968) that almost 7S'i



of atmospheric ammonia is removed from the atmosphere by com-eision to
NH, lonb which condense in cloud droplets or particles and may e\rapora('



to form aerosols,  Georgii (1963) found that in a polluted a.ir mass



containing approximately 500 ugSCL/m  and 10 ugNH_/m  , 98% of the anunon '



is converted to  (NH.)?SO. aerosols.  McConnell  (1973) has estimated tluj'



of a total NH  source strength of about 17,4 x 10   kg NFL,   N/yr,
             *J                                           -_)


approximately 3 x 10   kg of this ammonia is removed  each yeui by



rainout.  He points out that this is a very conservative estimate and



"may be low by as much as a factor of ,3."  Earlier, Robinson and Robbif



estimated that 280 x 10   kg NH,, - N are removed each year by precipita
                               O


tion.   On the whole, quantitative estimates are rather sparse and show


                                                                      4
a great discrepancy.  To amend  this, extensive measurements of the Nil



content of rainwater must be made and analyzed,,
     Atmospheric Reactions Involving NH  :  Until just recently the re^



involved in the destruction of atmospheric ammonia were poorly define-i



McConnell (1973) has suggested a number of reactions he believes to br



significant in its destruction.



     In the stratosphere the primary destructive reaction  involves the



photodissociation of NH  by sunlight with wavelengths greater than

     O

2300 A.  The basic reaction is
                                 hv -* NH  + H
-------
                                    -86-
                             -4    -1
The reaction rate is 1.3 x 10   sec  .   The various excited NH2 and NH



states possible, together with their observed threshold wavelengths and



their respective quantum yields are listed in McConnell's Table 1.



     Great quantities of hydroxyl radicals are now believed to exist



in the troposphere (Levy, 197]), the number density appearing to be on


                     12  -3
the order of about 10   m   (McConnell et al,, 1971).  Reaction of NH



with these OH radicals in the troposphere and lower stratosphere is



therefore possible, and, probably constitutes an important sink as well.



The reaction is as follows:





                          NH  + OH ->• NH  + HO




                                         -19  3    -1
Stuhl (1973) records the rate as 1.5 x 10    m  sec   at room temperature.



The magnitude of this rate constant is consistent with observations of



Worley, Coltharp and Potter (1972) and Albers, Hoyermann, Wagner and



Wolfrum (1969).  McConnell points out that this same reaction can and



probably does occur in the stratosphere, but it is not as effective



there as photolysis.  He further points out that although reactions of



NH, with 0, 0( D) and 0  are possible, evidence suggests that such



reactions would not play a significant role in the destruction of ammonia.



     According to McConnell, "the most likely atmospheric products"



of NH  destruction are N_ and NO.  In his paper he lists those reactions



most likely to occur along with any information that might be available



on their reaction rates.  An anomaly exists here that cannot as yet



be answered.  As McConnell indicated, oxidation of NH_ can be the  source
                                                     *!)


of a significant quantity of nitrogen oxides, especially NO  (estimated



at 7.04 x 10   kg N0y per year  as N).  But, he points out that the
-------
                                    -87-





magnitude of this NO source can be reduced due to a possible immediate



reaction of NO with NhL producing N  and HO.  This then would mean that



oxidation of NH  could constitute a sink for NO.  In effect then, oxida-
               o


ti'on of NH  can both produce and consume NO.  The proportion of NO



produced to that consumed by these reactions if they do occur is not



as yet clear, but his contention is that the NO produced far outweighs



that consumed.





Environmental Ammonia Cycle



     Relatively little quantitative information is available on the



strengths of both sources and sinks of ammonia.  In fact, research into



possible atmospheric oxidation reactions is so recent that any estimates



made are subject to gieat uncertainty.  Separate atmospheric ammonia cycles



have never been constructed; its cycle is always considered together with



that of nitrogen oxides.   Because not very much is even known about the



nitrogen oxides cycle, the uncertainty involved in the combined cycle



must be very high.



     To date the only nitrogen compound cycles devised are those of



Robinson and Robbins (1968, 1970) and McConnell (1973) and a total



geochemical nitrogen cycle of the present research,  The first three



cycles are compared in Table XV.  The discrepancy among the three in



source and sink estimates is immediately obvious.  More research must



be undertaken to provide us with a better understanding of the part



ammonia plays in atmospheric chemistry and to quantify possible sources



and sinks such as the oceans, lakes and vegetation.
-------
                                    -88-
Table XV.  ATMOSPHERIC FLUXES INVOLVED IN VARIOUS AMMONIA CYCLES
           (in 1010 kg N/yr)
Robinson Robinson
§ Robbins* § Robbins
(1968) (1970)
I. Sources
a. anthropogenic 	 .35
b. biological 670± 95.7
TOTAL 670 96.05
II. Sinks
a. precipitation 280 18.6
b. dry deposition 70 4.9
c. oxidation to NOY (troposphere) N.E. 	
A
d. oxidation § photolysis to NO N.E. 	
(stratosphere)
e, gaseous deposition 90 74.9
TOTAL 440 98.4
McConnell
(1973)

0.4
17
17.4

3
7
7
0.04
	
17.04
N.E. = Mechanism recognized but no estimate made.

*NOTE:  (see note in Table XI)

 ISource strength here was adjusted by Robinson and Robbins to
  provide an additional amount of nitrogen needed to balance other
  portions of their nitrogen compound cycle.
-------
                                    -89-
                              GLOBAL  NITROGEN  CYCLE

     The role which nitrogen compounds play in the chemistry of the
atmosphere and earth is not well understood.   A summary of what is known
and has been suggested about the routes and rates of transport of these
compounds is shown in Figure 4,
     The diagram has been divided into certain spheres or reservoirs in
which there exists an equilibrium concentration, or burden, of nitrogen.
These burdens have recently been compiled and, in some cases, calculated
by Wlotzka (1972).  The only elaboration made  upon his tabulations con-
cerns the abundance of the nitrite and ammonium ion present in the ocean.
Sverdrup, Johnson and Fleming (1942), as cited by Wlotzka, calculated the
NO" - N content of the ocean to be 5.7 x 10  Tg (5.7 x 10 7 g).  Emery,
  O
Orr and Rittenberg (1955)hkve estimated the total inorganic combined
nitrogen content, that is, N0~  N0~ and NH*  to be 5.8 x 10° Tg.   Therefore
in this cycle the N02 and NH  inventory of the hydrosphere has been set
equal to 0.1 x 10  Tg.  These burdens, shown encircled in their
                                                                 12
respective reservoirs in Figure 4, are given in units of Tg N (10  g N).
     There is so little quantitative data available concerning rates of
transfer, or fluxes, of these compounds between reservoirs it is impossible
to even attempt to balance this cycle.  As a result the cycle suggests
the possible routes of transfer and offers those fluxes which have been
                                               12
estimated.  Fluxes are expressed in Tg N/yr (10  g N/yr) and the direction
of transport is indicated by arrows.
     Details of the atmospheric portion of this cycle, with the exception
of biological nitrogen fixation and denitrification as a source and sink
of N  respectively, and the volcanic emanations of N , have been discussed
-------
                        -90-
2 z  NOUdaOSOV
   y  oooi - ooi

     Nouwxb 'oia
                                                                                   u
                                                                                   X
                                                                                  u


                                                                                   c;
 cd


 o
1—J



 CD


E-




«^-


 0)



 tifl
•H

tLi
-------
cocunc  removed CI*«'%T \,h;  .<  < •




bunion  :/.  NO., and ' r.1 ^  ':... ( ''-




and this  is  noted on the dia^i-jp.




      In formulating a  rutrogen bucifi. ;  •;  L i.  ,  . • ,•   i




aiso  estimated that 8  Tg N/yr j.-  at' '-. ••  • >< •  .t     :




fixation,  probably via some species  of  b im -^ieui aiv




An additional contribution to the t.-.-.,^  n -;  ..•, ,  -   ,




is surface and subsurface  2 and dru • j.if',!.




19 Tg N/yr are added to  the ocean .,}<  :,u-".'  '..;




(1972)  has caj.culatea  this co<'ur.i:), •  u.  ... i




      The most widely recognised prci « .   ,.i,.  ;




the ocean  is sedimentation   The <-'-r-  '  r,,   . L  •




manner  has been estimated  to range  • i". •  ,• ,/, />  •




8,6 Tg/yr  (Emery et al  ,  1955)   0. h- , -•:• •,...




include denitrification, producing .\   'i-J \,'.'
-------
                                   -92-
Liss and Slater (1974) calculated the flux of NLO from the sea to the air

at 38 ig/y.- (120 Tg N_0/>r).   Wlotzka estimated this flux to range between

20 and 130 Tg/yr.   The amount of N~ released from the ocean as a result

of denitrification is unknown, as is the amount of NH_ which escapes.

     Nothing is known about the contributions of mantle degassing,

volcanism, and weathering of the lithosphere to the total amount of

nitrogen carried to the ocean by land drainage.  It has been estimated

though that about 0.1 Tg/yr is released to the atmosphere by volcanic

activity.

     A great deal of nitrogen containing fertilizers are applied each

year to the soils across "he world.  Wlotzka recently estimated it to

total 30 Tg/yr.  Not only does this nitrogen play an important role in

the nutrient cycle between soil and plants, but much of it is subject

to leaching and would therefore contribute to the amount of nitrogen

removed by surface and subsurface land drainage, ultimately to wind up

in the sea,

     As mentioned previously, the quantity of nitrogen present in the

ocean is constant with tiir.e.  Based on the nitrogen budget of sea

water as formulated by Emery et al .  (1955), Dugdale  (1972) calculated

the residence time of nitrogen in the ocean to be:
       5
86 Tg/yr
                             ,.,0   9.2 x 10  Tg   In4
                         T = M'R =            * g 10
or about 10,000 years which is an extremely short period of time on the

geological time scale.  For the sake of comparison, it was earlier

calculated that the residence time of sulfate in the ocean is on the

order of millions of years.
-------
                                   -93-
     Figure 4 and the explanation above have attempted to show how vcix




little is known about the global nitrogen cycle.   Each question mark ii




the diagram needs to be replaced by a number representing an  actual I i




in the environment.   Those fluxes that have been  estimated need confine




tion.  Until all that is accomplished, any nitrogen cycle devised can -




be, at best, conjectural.
-------
                                   -94-
                                OXIDANTS





OZONE:



     The problems involved when significant amounts of ozone are present



in the atmosphere have come to light probably more as a result of the



photochemical pollution problem in Los Angeles than from any other single



factor.  While background concentrations of ozone probably range from



about 20-60 yg/m  (0.01-0.03 ppm), in urban centers like Los Angeles,



it is not unusual to have 0  present at levels greater than 500 yg/m
                           O


(.25 ppm).  Ozone levels up to 400 yg/m  (0,2 ppm) usually will not cause



any deleterious effects  (Masters, 1971), but at concentrations of 600 yg/m



(0=3 ppm) ozone causes irritation of the mucous membranes in the nose



and throat.  At. somewha- higher levels, it can cause coughing, choking



and severe fatigue.  When present at relatively high levels, such as



those that occur in severe photochemical smogs, ozone causes bronchial



irritation and interferes with normal lung functioning, causing breathing



difficulty and chest paini.  The highest ozone concentration detected in



the Los Angeles atmosphere was 2,000 yg/m  (0.99 ppm) in 1956  (Chambers,



as cited in Tebbens, 1968).



     Another serious problem related to the occurrence of atmospheric



ozone  is its toxic effect on vegetation, especially field and  forage



crops  (such as tobacco), ^.eafy vegetables, shrubs, fruit and forest trees



(particularly conifers).





Sources



     The maximum ozone density occurs at an altitude of 25(±5) km, in



the mid-stratosphere.  This ozone though is not formed there but between



30 and 60 km where the following reaction proceeds:
-------
                                    -95-
                          0+0 + M + O+M
                           £.            O




In this altitude range though ozone is relatively unstable.  It may be



destroyed either by collision with atomic oxygen to form molecular oxygen



or by photolysis.  The cycle is continuous; formation, destruction,



formation. .  ,    As a result an approximate state of equilibrium exists



above about 40 km.  The accumulation of ozone at 25 km is a result of i'<



downward transport to a location where its destruction is less likely



(Barry and Chorley, 1970).



     Because the wavelengths of ultraviolet radiation that penetrates



the troposphere are too long to cause photodissociation of oxygen, it



has long been accepted that the presence of ozone in the troposphere is



due primarily to the transport of ozone down from the stratosphere.  Th ;



amount present in the troposphere would then be related to the injectun



rate through the tropopause,  estimated as ranging between 1.9 and 7.5



x 10"5 kg/yr (Junge, 1962).



     Recently,  Chameides and Walker (1973) proposed a model wherein both



seasonal and diurnal variations in the tropospheric ozone density are



assumed to be caused by photochemical changes rather than a change in



the flux of stratospheric ozone-rich air into the troposphere.



     Their model calls for the production of 0  by the methane oxidation
                                              J


scheme suggested by Crutzen (1973) .   The oxidation of methane produces



hydroperoxyl radicals which then react with nitric oxide:





                      H02 + NO ->• OH + N02 followed by                (} >





                      NO  + hv -* NO + 0 and                          (/ i
-------
                                  -96-
                      0 + 02 + M + 03 + M                            C3)





At 27°C, k  is about 5 x 10"   m  sec"  (Levy,  1973a).   Based on calculated



atmospheric densities of HO  and NO , Chameides and Walker calculated an


                                             12  -3    -1
ozone production rate of approximately 5 x 10   m   sec  , and photo-



chemical lifetimes of a few tenths of a day near the ground,  of one day



at 5 km, and 10 days at 10 km.  They concluded  that ozone  is  in a state



of photochemical equilibrium in the troposphere.



     Upon developing a photochemical model of reactions affecting the



density of tropospheric ozone, they found they  could satisfactorily



reproduce the seasonal variations found in ambient ozone densities (see



Figure 5) by applying their model to temperature and relative humidity



conditions typical of winter and summer.  By analyzing the specific



effects that lowering the temperature, humidity and photodissociation



rates each make in altering the summer profile  to one characteristic of



the winter, they conclude that decreasing the photodissociation rates



alone causes a realistic alteration of the summer profile to  that of



winter  (Figure 6).  This model shows that photochemical processes rather



than the vertical transport of 0 -rich air can be responsible for the
                                o


presence of ozone in the troposphere.  It also presents a challenge for



other researchers to prove whether or not most  tropospheric ozone



originates in the stratosphere as commonly believed or in the troposphere



by the  reactions these authors propose.





Removal Mechanisms



     Ozone is relatively insoluble in water; at 20°C and  1 atm its solubi-



lity is 0.052 g/lOOg HO.  Therefore, removal of ozone from the atmosphere
-------
                                           -97-

^ ^^
^ J
^^ — •
^^
^s»
\^_^^ ^ — ^

""" — 	 ••• ==^sS9ra5.:=

— . — _. 	 	

„
i . 1 . i , 1 . 1
-••*. *^\ a f\ 	 i_ f± 1 >••*.
OJ

O

CD

CD

_^L
 | 1
4_j :P
t' r'
4) ">
tw fj
f . , rt . ,i
H-4 Xs
0 H i'
Pi '.'')
1.. f. ;»
xr, <•< i
!." fJ
'!>
f-t
3
b.0
.-J

photodissociation rates
decreased to their winte:
values

i
!
I










          00
          CD
CVJ
<3
 O
                                                          00
                 to

                   o

                   CO
                                                               o
                                                          CVJ
                                                                fO
                                                              O
co
                      CD
CVJ
o
                                                                     C'
                                                                     p.
                                                                     13
                                                                     QJ  f '

                                                                     I"
                                                                                       10
                                                                                       a)
                                                                                      •H
                                                                                        	 /
•H
•4-1 ^-4
• H  nj
C
•H
f-l
X

X
•^
T)
CD
(H
3
in
0) TJ-
6 o
en
co i— i
CD ' — '
•r-t
•*-* P«
•H a>
CO T)
r1 J-*
0) O
-------
                                   -98-





by washout and rainout can be disregarded.   There is evidence though that



ozone is removed by the oceans.   Ozone is also absorbed by vegetation and



soil and these appear  to  represent  major  sinks,  for  0  .   Due  to  its
                                                    o


nature as a strong oxidising agent, ozone participates in a number of



atmospheric reactions, especially in polluted atmospheres.  Junge (1962)



has estimated the tropospheric residence time of ozone to range between



3 and 4 months.





      Vegetation:  Ozone has caused a great deal of damage to vegetation



in the United States and its effects on both plants and the economy have



been well documented  (Rich, 1964 and Millecan, 1971).   When ozone is



present in the atmosphere at low concentrations, vegetation is capable of



absorbing some without inflicting damage upon itself.   Hill (1971) showed



that an alfalfa canopy removed ozone from his experimental chamber at a



rate of 1.7 cm/sec.   In this study the threshold level above which the


                                     -43          3
plants sustained injury was 4.16 x 10   mol/m   (200 yg/m  or 0.1 ppm).



Previously, Hill and  Littlefield (1969) had shown that high concentrations



of ozone can cause stomatal closure in the leaf thereby reducing the



uptake of ozone.  Unfortunately in Hill's study the injury threshold limit


                        -4      3
of alfalfa was 4.16 x 10   mol/m  which was also the level above which the



stomates began to close.  As a result, the effect of stomatal closure on



ozone uptake rates could not be adequately evaluated.



      The hypothesis  that the uptake of 0  is dependent on the degree of
                                         O


stomatal opening received additional support by Rich, Waggoner and



Tom1inson  (1970) as a result of their experiments on bean plants.  They



concluded that ozone  follows the same but reverse path as that of water



vapor which is transpired by the leaf.  The water vapor originates from
-------
                                   -99-
saturated cells directly beneath the stomata and is released when the



stomata are open.  Conversely, the ozone travels through the open



stomata onto the surface of the substomatal cells where it is reduced to



very low concentrations.  Hill points out that although ozone's solubility



in water is not very high, it breaks down relatively rapidly and therefore



a fairly high uptake rate is to be expected.





      Soil and Water Bodies:  The ground has long been recognized as a



sink for ozone.  Over the last two decades several estimates of the



vertical flux of 0  into the ground have been made.  Junge (1962) tabulated



the estimates made through 1962 [see his Table I].  For the most part,


                                            -12       2
the estimates range from about 3 to 200 x 10   kg 0 /m -sec.  Jurige calcu-


                                    -12     2
lated the flux density to be 60 x 10    kg/m *sec which is in good agreement


                            -12     2
with the estimate of 48 x 10    kg/m 'sec made by Kroening and Ney (1962).



These measurements of 0  into the ground include destruction by both soil



and vegetation.  No attempts were made to differentiate the ozone destroyed



by each, and no suggestions were offered as to what the exact reason for



the destruction at these surfaces were.



      Aldaz  (1969)  was  the first  to  offer data  on the  ozone  flux  into  various



surfaces: soil, vegetation, snow and water  (Figure 7).  These fluxes were



determined assuming a partial ozone density of 40 yg/m  (0,02 ppm, or



natural background concentration levels).  Using an equation analogous



to that used by Liss and Slater (1974) to estimate the flux of a gas



across the air-sea interface, Aldaz calculated the flux of 0  into both



the terrestrial and marine portions of the northern and southern



hemispheres.  The equation used to determine the ozone flux is





                            F = 1.25 x 1010 kq
-------
                       -100-
  10
    16
 o
 0>


CM



 l|015
 to
o
O
X
o>
c
o

5
  10'
  10
    13
       ^JUNIPER BUSH (N.M.)



       ^-TUNDRA(ALASKA) KELLEY, 1968.


       — SAND OR DRY GRASS (N.M.)

       ^- GRASS (AUSTRALIA) GALBALLY, 1968.

       — GRASS (NEBRASKA) REGENER, 1957.


       — SNOW (N.M.)
          WATER(FRESH)(N.M.)


          WATER(SEA)(40°N, 70°W)

          WATER (SEAKTROPICAL, MIDDLE LAT.)
         1VATER (DISTILLED)
          CLEAN MYLAR
     Figure 7. Ozone fluxes; into different surfaces assuming a partial

            ozone density of 40 yg/m^.
-------
                                   -101-
where q is the partial density of 0   in yg/m  ,  and  k,  the j" •  i



is a measure of the chemical reactivity of  the  surface to -^i',"



then, k is actually a deposition velocity term.   The  react !••>



used were 0.60 cm/sec over land, 0.04  cm/sec  over water aii"  '



snow.  Data Aldaz used to compute the  global  ozone  sink i s :.- •



XVI.  Because there are no measurements of  ozone  uptake by > •  i



tion, Aldaz assumes in estimate  (A) that the  rate of  destriu . •



tropics is equal to that in other land regions, and in (BJ ;,



destruction rate 5 times greater in the tropics than  in ot!:<



In conclusion, he found the sink strength of  the  earth's si',1 -•



between 1.3 and 2.1 x 10   kg/yr.  He  also  showed that de';l<'



ozone by New Mexican soils is about 15 times  faster than ».  -,  •



Ocean (Figure 7), and that bare, dry  soil destroyed approxi'-1



more 0  than when moist   It must be pointed  out  that  the =, -. :



which the soil destroys ozone was not  identified.



      Turner, Rich and Waggoner  (1975) recently found thcil   • .•



freshly cultivated, fine sand loam removed  ozone  at a iat._ -'••:


           -8      2
.5 - 2 x 10   mol/m -sec when present  in the  atmosphere uf



ranging from 64-208 ng/m .  They also  found that  as the soi i :,.



content was increased, the resistance  of the  soil to  0,, !••"• i . i



increased.  As mentioned above, Aldaz  found a similar dir.i...



in the sand loam soil he tested in New Mexico;  i.e.,  as li<



increases, ozone destruction by the soil decreases,
      Atmospheric Reactions  Involving 0  :   Although  react.-.
with natural contaminants such as terpene  do  occur



1965; and Ripperton et al . , 1967),  its  importance arises  I
-------
-102-



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-------
                                     -103-
                      *
photochemical reactions in which  it  participates  in poll-"'

namely the photo-oxidation of hydrocarbons  in the preben ,

dioxide.  The initial oxidation of olefins  by ozone foi  *<•

to a long series of reactions which  produce ketones,  al-f1'

acids and nitrogen-containing compounds  such as peroxy;s  ••>

The presence of these compounds in the atmosphere has !•• -."•

cause of considerable eye irritation (Schuck and  Doyle,  i

constants for the reaction of ozone  with numerous hydr-<  ••

summarized by Altshuller and Bufalini  (1971)  and  BufaJ <  '

(1965).  Because the reactions of ozone  in  polluted a1 ir   <

numerous and have been the subject of extensive invest i;; '•

review, it would not serve the best  interest of this  p::  ••

reactions here.   For additional information the readei

Hecht and Seinfeld (1973), Dutsch (1971), Ripperton an,!   ,

Stepnens (1969), Altshuller and Bufalini (1965),  and  I...  • •
-------
                                    -104-
                             ORGANIC GASES






      Organic gases constitute the second major group of air contaminants.




This group includes all classes of hydrocarbons and those formed when some




of the hydrogen of the original compound is replaced by other substituent




groups containing nitrogen, sulfur, or oxygen.   Organic gases are further




subdivided into reactive and non-reactive classes.   Among these classes




only hydrocarbons are cor side red here.  The more important reactive




hydrocarbons include the olefins and aromatics.  Paraffinic hydrocarbons




are classified as non-reactive,






REACTIVE HYDROCARBONS




      Because of their role in photochemical reactions in polluted




atmospheres, reactive hydrocarbons have been the subject of great interest.




These photochemical reactions produce smog which is associated with eye




and respiratory tract irritation, reduced atmospheric visibility, and




plant damage.




      Limited ambient data from Point Barrow, Alaska  (Robinson and Robbins,




1968)indicate that ethylene, the most abundant hydrocarbon of this group,




is present at a concentration of less than 1 ug/m   (less than 1 ppb).




In the absence of other ambient measurements one might consider the




concentration of ethylene measured at the above location to represent




the upper limit of the background concentration for components in this




group„






Sources




      A variety of hydrocarbons are released to the atmosphere as a result




of both anthropogenic activity and natural processes.  The most important
-------
                                    -105-
anthropogenic source of hydrocarbons is motor vehicle exit,:.'


incomplete combustion of fuel.  Robinson and Robbins  (llJbSi


that the total annual emission of olefins and aromatics  tV;

                                9
tion of various fuels is 27 x 10  kilograms.  A bibliograj.i


emission sources has been compiled by the U.S. Department  <•>


Education and Welfare (1970).


      Plant species also release appreciable quantities  or


organic substances to the surrounding air.  The major rc.i  • •


emitted by trees are ethylene, monoterpene  (C,n), and isop,


a recent study, Rasmussen (1972) concluded that the
                                   9
a global natural source of 175 x 10  kg of reactive hydiv*


year.   This emission rate is 6 times greater than that csi . •


reactive hydrocarbons of anthropogenic origin,



Removal Mechanisms


      Hydrocarbons in general are not water soluble, and  ' !


cannot be directly removed from the atmosphere by wet pix» >


washout and absorption by surface waters,  Various studios '


that photochemical reactions are important in removing roa'


carbons, although the products formed may cause detriment.1 '


such as eye and throat irritation.  This class of hydroca.i'


their emission into the atmosphere undergo rapid chemical  >'


in the presence of: atomic oxygen and ozone (Bufalini and  '


1965); oxides of nitrogen (Schuck, 1961; Alley, Martin and


ozone and sulfur dioxide (Cox and Penkett, 1971b) and; n.i i


and sulfur dioxide (Schuck and Doyle, 1959)
-------
                                   -106-






      The basic kinetic mechanisms of hydrocarbon reactions in the



atmosphere are given by Hecht and Seinfeld (1972).   These authors present



a 15-step mechanism for photochemical smog formation, with rate constants



and stoichiometric coefficients chosen according to the particular hydro-



carbons involved in the reactions and the initial reactant ratios.  The



state of the art of photochemical reactions is analyzed by Dodge  (1973)



and Seinfeld, Hecht and Roi:h (1973).   The reader is referred to these



publications for an up-to-date view.



      Quantitative data on the rates of these atmospheric reactions are



rare.  The limited data available as reported by Hidy (1973) suggests



that 1-10% by weight of the; reactive hydrocarbons emitted into the



atmosphere are converted to aerosols and eventually removed by scavenging



or deposition.  The remaining hydrocarbons are eventually oxidized to



carbon dioxide and water vapor.



      Smith et al.  (1973) investigated the capacity of soils to absorb



ethylene and acetylene.  Sterilized and air-dried unsterilized soils



removed neither of these hydrocarbons from the test chamber.  Moist



unsterilized soils  (50% saturated) did reduce the ambient concentrations



of these gases although the uptake rate was relatively slow compared



to the other gases they studied  (S0_ and CO).  They concluded, as did



Abeles et al.  (1971) in an experiment on ethylene uptake by soil, that



the  sorption of both ethylene and acetylene is due to microbial activity



in the soil.   Smith et al. found that the soils they tested removed


                                                   -9
ethylene at average rates ranging from  .14-.97 x 10   mol per gram of


                                                       -9
soil per day  (mol/g*d) and acetylene from  .24-3.12 x 10   mol/g*d.
-------
                                    -107-
NON-RL ACT[VL IIYDROCARBONS


      This group, which  consists  of methane and the higher saturated


hydrocarbons, has been found to be  much less involved in photocliunii1,, i


reactions and smog formation than reactive hydrocarbons.  By jar ! lir  •


abundant paraffinic hydrocarbon in  the  atmosphere is methane,  i'letb;-h:


background concentration is about 1000  yg/m  (1.5 ppm) ,  while Lh,/


background concentration for heavier gases in this class is ie: ,s I h,n)


1 yg/m" (less than 1 ppb)  (Robinson and Robbins, 1968),



Sources


      The anthropogenic  source of paraffinic hydrocarbons is the in-i.,


combustion of fuel in motor vehicles.   The annual emission o! i-,ira!i i<

                                        9
hydrocarbons is estimated  to be 60  x 10  kg (Robinson and Kolihui:,   i

                                                   9
      Among natural sources approximately 310 x 10  kg of medi-iiie i.-.


produced annually in swamps and various water bodies as  a revtlt  «i


bacterial decomposition.   The relatively high concentration of moth.,'.


in the atmosphere compared to other organic gases is related to tin--.


natural process.



Removal Mechanisms


      Because they are so  insoluble in  water, paraffinic hydroea i }jf>.<


can not be removed from  the atmosphere  by wet processes.  The J.H'IJVJ; ,


sink for methane in the  troposphere is  its oxidation by  hydroxyl  i;),;:•


to form carbon monoxide.   The initial reaction is as follows:
                                CH   +  OH -> CH  + HO
                                  •           O     *-
-------
                                  -108-




                                                      -12
The rate coefficient for this equation equals 5.5 x 10    exp (-1900/T)



(Grciner, i970) .   For a complete development of this oxidation scheme the



reader is referred to Levy (1971), McConnell et al. (1971) and Levy



(1972, 1973a).   Based upon density profiles of CH  and hydroxyl radicals



in the troposphere, and the rate equation given above, Levy (1973b)



calculated the average daily loss of methane at a particular altitude.

                                                                _ Q

The total column loss rale for methane was found to be 7.48 x 10


     2
mol/m -sec.  This results in a tropospheric residence time for methane



of 2 years.



      Rasmussen,  Hutton and Garner (1968) and Robinson and Robbins (1968)



have also suggested that the volatile organic components of the atmosphere



are removed by bacteriological processes and vegetation.  However,



quantitative data and rate equations for these removal mechanisms are



nonexistant at present.
-------
                                    -109-
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                                   -110-
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-------
                                    -Ill-
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                                    -112-
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Wickman, F, E,  personal communication, May 8, 1974,

Wlotzka, F,  "Nitrogen cycle" in Handbook of Geochemistry, Vol. H/3,
     K. II, Wedepohl, ed,, Springer-Verlag, Berlin, pp,  7-0-1 to 7-0-3,
     1972.

'A'oriey, S. 0. , R. N  Coltharp and A- E, Potter-  "Rates of  interaction of
     vibrationally excited hydroxyl (v-91 with diatomic and small
     polyatomic molecules,"  J.  JPhys  Chem., 76:1511-1514  (1972).
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                                   TECHNICAL REPORT DATA
                            >'L jw r':t;JI>-i"jj-,tn on the rc;'t'rse before complsti'ig)
                                                          J3, RECIPIENT'S ACCESSION-NO.
            u; ;Hd".;,,r.'/.  Removal Processes  for
             p'ieric rollutants
             Kabel
                          Mamoo?1  Faheri
                          (Pahlavi Univ.,  IRAN)
*v:,NG OH "ANIMATION NAMb AND ADDRI-.JS
•nr  for Air Environment Studies
renske Laboratory
Pennsylvania State Uni versitv
•ersity °ark. PA  16802
J5. REPORT DATE
j_	
6 PE.RF-ORMING ORGANIZATION CODE


8 PfcHI-OHMING ORGANIZATION HI COM I MO

  CAES Publication No. 367-74
     •• TOfiiNG AGENCY NAMfc AND ADDRbSS
     ;eorology  Laboratory, EPA
     /onal  Environmental Research  Center
     ectrch  Triangle Park, N. C.  27711
                                                           10 PPCiGRAM ELEMENT NO.
                                                                     1A1009
                                                           11. CONTRACT/GRANT NO.

                                                                     800397
                                                      K3. TYfp OF REPORT AND PERIOD COVERED
                                                       	I terim	
                                                    j 14. SI?ONS^OR.TN"G AG₯^Jcv CODE
               MOTES

••I. A3STRACT

  T.his  review attempts to briefly  "llustrate what the  "state of the art"  is  in  the
  recognition of the various soum;s  and natural sinks  of gaseous pollutants.   The
  removal  mechanisms include absorption by vegetation,  soil, rock and water  bodies,
  precipitation scavenging, and  chemical reactions within the atmosphere.  The
  nature and  magnitude of anthropogenic and natural emissions of the gases
  considered  i%S, SCU NO, NO, NO , NH0, CO, 0 , and  hydrocarbons), along  with
  r,revr ambient^backgrouna concentrations' and information on their major  sinks
  •identified  to date, are discussed.   In the case of sulfurous and nitrogenous
  cor.iooundSj  this information  r-.as  neen used to prepared total geochemical  cycles.
                  DESCRIPTORS
                         KEY WORDS AISiD DOCUMENT ANALYSIS

                                       jb. IDENTIFIERS/OPEN ENDED TERMS
                  Hydrogen sulfide
                  Sulfur dioxide
  ;.".  reactions    Carbon monoide
 • ;".cr=             NitrotiS oxide
                  Nitric oxide
                  Nitrogen dioxide
                  Ammonia, Hydrocarbon
                                                       Sinks
                                                       Sources
                                                       Natural  removal
           i:  CVClSS
           :yc! ~
                 uzone
          O.\ STATEMENT
             nn ted
                                             ,19. SECURITY CLASS (This Report)
                                             ;      Unclassified
                                             J20. SECURITY CLASS (This page)
                                             \      Unclassified
              c.  COSATi 1'iclci/Group
                                                                         21. NO OF PAGES
                                                                   22
                                           122
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